2.-Atohjejrjjejekekekekemic-Structure.pptx

1.
Atomic
Structure
Atomic Structure

2.
Atomic
Structure
6tom
⚫ Each element is composed of
smallest particles known as ‘ATOM’.
⚫ Atom the name is derived from Greek
word Atom means - ‘Not to be cut’.
D6LTON’S THEORY OF 6TOM
J. Dalton developed his famous theory
of atom in 1803. The main postulates of
⚫ Atom was considered as a samllest,
dense and hard indivisible particle.
⚫ Each element consists of a specific type
of
atoms.
⚫ The properties of elements differ because
of difference in kinds of atom contained in
them.
⚫ This theory provides a satisfactory basis
for law of chemical combinaiton.
Limitations of Dalton’s Theory
⚫ Dalton’s theory fails to explain why atoms
of different kinds should differ in valency
and mass etc.
⚫ The discovery of isobars and isotopes
showed that atoms of same elements have
different atomic masses (isotopes) and
atoms of different kinds may have same
atomic masses (isobars).
Concept Ladder
Atoms is
indestructible, i.e., it
cannot be created or
destroyed.
Racm your Brain
Why and how do atoms
combine together to form
compound atoms?
Definition
Each element is composed
of smallest particles called
‘ATOM’.
Concept Ladder
Democritus was the
person who first
suggested the existence
of ATOM & coined the
name ATOMOS means Not
to be cut or Indivisible.

3.
Atomic
Structure
⚫ The discovery of various sub-aomic
particles like protons, electrons, X-rays etc.
during late 19th century lead to idea that
the atom was no longer an indivisible and
smallest particle of the matter.
DISCOVERY OF FUND6MENT6L P6RTICLES
⚫ Atom consist of several sub-atomic
particles like neutron, proton, electron,
neutrino, positron etc. Out of these
particles, electron, proton and the neutron
are called fundamental subatomic particles.
(1) Electron
⚫ Electron discovered by J.J. Thomson (1897)
and it is negatively charged particle.
Cathode Ray Experiment
⚫ William Crookes in 1879 studied the
electrical discharge in partially evacuated
tubes called as cathode ray discharge tube.
Concept Ladder
R.S. Mullikan measured
the charge on an
electron by oil drop
experiment. The charge
on each electron is –1.602
× 10–19
C.
Racm your Brain
Why, cathode rays do not
depend upon the nature of
gas or the cathode meterial
used in discharge tube?
⚫ Discharge tube is made of glass, about
60 cm long containing two thin pieces of
metals known as electrodes, sealed in it.
This is called crooke’s tube. Negative
electrode is called as cathode and
positive electrode is called anode.

4.
Atomic
Structure
⚫ When a gas enclosed at low pressure (~10–
4
atm) in discharge tube is subjected to a
high voltage (~10,000 V), invisible rays
originating from the cathode and
producing a greenish glow behind the
perforated anode on the glass wall coated
with phosphorescent material ZnS is
observed. These rays
Properties
⚫ Cathode rays travel in straight line.
⚫ Cathode rays produce mechanical effect,
as they can rotate the wheel that placed in
their path.
⚫ Cathode rays consist negatively
charged particles called as electron.
⚫ Cathode rays travel with high speed.
⚫ Cathode rays can cause fluorescence.
⚫ Cathoderays heat the object on which
they fall due to transfer of kinetic energy
to the object.
⚫ When cathode rays fall on heavy
metals, X-rays produced.
⚫ Cathode rays possess ionizing power that
is they ionize the gas through which they
pass.
⚫ The cathode rays produce scintillation on
the photographic plates.
⚫ They can penetrate through thin
metallic sheets.
Racm your Brain
Why do cathode rays
produce
fluorescence?
Concept Ladder
Order of specific
charge
 n  p 
e
 e    e  
 e 
 m
  m
 
m 
e –
 mass of proton 

mp

1837
 mass of electron m


Previous Year’s uestion
Ǫ
Cathode rays have
[6IPMT]
(1) Mass only
(2) Charge only
(3) No mass and charge
(4) Mass and charge
both

5.
Atomic
Structure
DISCOVERY OF FUND6MENT6L P6RTICLES
Charge by mass ratio of Cathode Ray
Electron is a low-mass, negatively
charged particle. It can easily be swing by
passing close to other electron or positive
nucleus of an atom. With the help of
variation in electric and magnetic fields the
ratio was determined. The
apparatus is shown below.
The charge to mass ratio of the electron is
given
Charge to mass ratio of electron is given
by: e/m = 1.758820 × 1011
C/kg
Where;
m = mass of electron in kg.
e = magnitude of the charge of electron
in coulombs.
(2) Proton
⚫ Proton was discovered by Goldstein and it
is positively charged particle.
6node Ray Experiment
Canal Ray experiment is the
experiment performed byGerman
scientistEugen Goldstein in 1886 that led to
the discovery of the proton. The discovery
of proton which happened after the
discovery of the electron further
strengthened the structure of the atom.
In the experiment, Goldstein applied
high voltage across a discharge tube which
had a perforated cathode. A faint luminous
ray was
Concept Ladder
Anode rays, e/m value
is dependent upon the
nature of the gas taken
in the tube. It is
maximum when gas
present in the tube is
hydrogen.
Racm your Brain
What were the important
conditions maintained in the
discharge tube by Goldstein?

6.
Atomic
Structure
seen extending from the holes in the back
of the cathode.
Properties
⚫ Anode rays travel in wtraight line.
⚫ Anode rays are material particle.
⚫ Anode rays are positively charged.
⚫ Anode rays may get swing by external
magnetic
field.
⚫ Anode rays also affect the photographic
plate.
⚫ e/m ratio of anode rays is lesser
than electrons.
⚫ These anode rays produce flashes of light
on
ZnS screen.
(3) Neutron
⚫ In 1932, Chadwick bombarded Be with
a stream of -particles (He+2
). He noticed
that penetrating radiations were produced
which not affected by magnetic field and
electric field. These radiationss consisted
of neutral particles, which were called
neutrons.
Ǫ
The discovery of neutron
become very late because
[6IPMT]
(1) Neutrons are present in nucleus
(2)Neutrons are highly
unstable particles
(3) Neutraons are chargeless
(4) Neutrons do not move
Concept Ladder
Anode rays were observed
by E. Goldstein but
was named by E.
Rutherford.

7.
Atomic
Structure
4 2
Be9
+ He4  6 0
C12 + n1
5B11 + 2He4  7 0
N14 + n1
⚫ These radiation was made incident on
paraffin wax, a hydrocarbon having a
relatively high H2
content.
⚫ These protons ejected from the paraffin
wax (when struck by uncharged radiation)
were observed with the help of ionization
chamber.
⚫ The range of the liberated protons
was measured and the interaction
between the uncharged radiation and the
atoms of several gases was studied by
Chadwick.
⚫ The neutron is relatively massive but
neutral, it is scarcely affected by the cloud
of electrons surrounding the nucleus or
by the positive electrical barrier of the
nucleus it self, thus it can penetrate the
nucleus of any element
Concept Ladder
Neutron is
fundamental particle of
all the atomic nucleus,
except hydrogen or
protium.

8.
Atomic
Structure
Particles Symbol Mass Charge Discovered By
Electron –1
e0
or 
9.1096 × 10–31
kg
0.000548 amu
–1.602 × 10–19
Coulombs
–4.803 ×
10–10
esu
J.J.
Thompson
Stoney
Lorentz
1887
Proton
1
H1
1.6726 × 10–27
kg
1.00757 amu
+1.602 × 10–19
Coulombs
+ 4.803 × 10–10
esu
Goldstein
Rutherfor
d 1907
Neutron
0
n1
1.6749 × 10–27
kg
1.00893 amu
1 amu  1.66 ×
10–27
kg
Neutra
l 0
James
Chadwick
1932
MODELS OF 6TOM
(1) Thomson’s Model
⚫ Thomson was the first proposed a
detailed
model of atom.
⚫ Thomson proposed that an atom consist
of a uniform sphere of positive (+ve)
charge in which electrons are present at
some places.
⚫ This model of atom is known as
‘Plum- Pudding model’.
Concept Ladder
The colloquial nickname
‘plum pudding’ was soon
attributed to Thomson’s
model as distribution
of electrons
withi
n
region
many
its positively
charged of space
reminded scientists
of raisins,
then called ‘plums’.
Racm your Brain
What is the name of
Thomson’s model?

9.
Atomic
Structure
Limitations of Thomson’s Theory
⚫ An importnat drawback of this model is
that the mas of the atoms is considered
to be evenly spread over that atom.
⚫ It is a static model. It does not reflect
the
movement of electron.
(2) Rutherford’s -Scattering Experiment
⚫ Rutherford carried out -particles
scattering experiment by the bombardment
of high speed
-particle on thin foil of gold, emitted
from radius and gave the following
observations, which was based on his
experiment.
⚫ Theangular deflections of
scattered
-particles were studied with the help
of moving microscope.
Concept Ladder
The relation between
number of deflected
particles and deflection
angle  is
 
1
sin4

2
Where,  = deflected
particles
 = deflection
angle
Ǫ
The nucleus of atom consists of
[6IPMT]
(1) Proton and neutron
(2) Proton and electron
(3) Neutron and electron
(4) Proton, neturon and electron

10.
Atomic
Structure
Observations
Rutherford execute a number of
experiments, involving the scattering of -
particles (He+2
) by very thin foil of gold.
Observations were:
⚫ Most of the -particles (99%) passes
through
it, without any deviation or deflection.
⚫ Some of the -particles were
deflected
through small angles.
⚫ Very few -particles were deflected by
large angles and occasionally an a-
particle got deflected by 180°.
Conclusions
⚫ Most of the -particle (He+2) went
straight through metal foil undeflected, it
means that there must be very large
empty space within atom.
⚫ Since few of the -particles were
defelcted from their original path
through moderate angle; it was
concluded that whole of the positive
(+ve) charge is concentrated and the space
occupied by this positive charge is very
small in the atom.
⚫ When -particles comes closer to this
point, they bear a force of repulsion and
diverge from their path.
Concept Ladder
An atom consists
of positively charge
center which is known as
nucleus e–
revolves
around nucleus (+) like
solar system
Ǫ
Rutherford experiment prove
[6IPMT]
(1) Electron
(3) Atom
(2) Proton
(4) Nucleus

11.
Atomic
Structure
⚫ Positively (+ve) charged heavy mass
which occupies only a small volume in an
atom is known as nucleus. Nucleus is
supposed to be present at the centre of
atom.
⚫ A very few of the -particles (He+2
)
suffered strong diversion on even
returned on their path indicating that
nucleus is rigid and
+2
-particles (He ) recoil due to direct
collision
with heavy positively charged mass.
⚫ As atomic number increases, the
number of protons increases which
increases the
repulstuion and so deflection angle

increases.
⚫ Atom has two
part
(1) Nuclear
part

Mass
(2) Extra nuclear
part

Size, volume
Size of nucleus = 10–13
cm or 10–15
m = 1
fermi Size of atom = 10–8
cm or 10–10
m = 1
Å
= N
 N

–5

10
A A
r
D
10–13
r
D
10–8
⚫ Density of nucleus = + 1017
kg /m3
Limitations of Rutherford’s -
scattering experiment
⚫ It does not obey Maxwell theory
of electrodynamics. Charge particle in
attractive field revolves to emits its energy.
So It’s loop is reduced and collide with
nucleus.
Concept Ladder
⚫ Mass Number (A) =
Number of Protons (Z) +
Number of Neutrons (n)
⚫ Number of Neutrons (n)
= Mass Number (A) –
Number of Protons (Z)
Concept Ladder
1/3
n
R  A
A = mass no
Rn
= R0
{A} 1/3



R0  1.33 fermi
  1.33  10–13
cm
 
3 3
n 0
R  R A
Rn
= Radius of nucleus
0
R = constant
model
Racm your Brain
Why was
Rutherford’s rejected?

12.
Atomic
Structure
⚫
But in reality discontinuous spectrum
occur.
6TOMIC NUMBER 6ND M6SS NUMBER
(1) 6tomic Number (Z)
⚫ It is always equal to number of protons
present in the nucleus of the atom of the
element. It also represents the number
of electrons in the neutral atom.
Ex: Number of protons in K is 19, so its
atomic number (Z) = 19.
(2) Mass Number (6)
⚫ The neutrons and protons present in
nucleus of an atom are collectively known
as nucleons.
⚫ An element have two isotopes Z1
and
Z2
, that of isotopes are W1
and W2
and their
% of occurance in nature are X1
and X2
respectively then the average atomic
weight of element is —
avg. wt 
W1
X1
 W2
X2
X1
 X2
Ǫ
The number of electrons and
neurons of an element is 18 and
20 respectively. Its mass number
is :
[6IPMT]
(1)
17
(3) 2
(2) 37
(4) 38
Concept Ladder
Atomic weight may
be decimal but mass
number of atom always
a whole number.
Racm your Brain
How
do
you
find
the
atomic
number?
6tomic Weight
⚫ It is the average of average of weights of
the all isotopes of that element.

13.
Atomic
Structure
ISOTOPES, ISOB6RS, ISODIPHERS,
6ND ISOELECTRONIC
(1) Isotopes
⚫ Isotopes given by Soddy.
ISOSTERS
⚫ Different atoms of same elements which
have similar atomic number but
different mass number.
⚫ The chemical properties are controlled
by the number of electrons. Thus, isotopes
of an element show same chemical
behaviour but different physical properties.
Ex: 1
H1
, 1
H2
, 1
H3
17
Cl35
, 17
Cl37
6
C12, 6
C13, 6
C14
8
O16, 8
O17, 8
O18
1
H1
1
H2
1
H3
e – 1 1 1
p+ 1 1 1
n 0 1 2
*Similar e–
, P+
different n
*n + p = 1,2,3
(nucleons)
different no. of nucleons
*Nuclear charge same
Racm your Brain
Isotopes and Isobars both do
not have same value of e/m.
Why?
Concept Ladder
There are mainly
types of isotopes.
are radioactive
and
tw
o
These
stable
isotopes. Stable
isotopes have a stable
nuclei and do not undergo
decay.

14.
Atomic
Structure
(2) Isobars
⚫ Isotopes given by Aston.
⚫ These are the different atoms of
different elements which have similar
mass number but different atomic number.
⚫ Isobars do not show the same
chemical properties.
Ex: 6
C14
, 7
N14
17
Cl37
, 18
Ar37
3
Ce76
,
2 34Se74
2 6
Fe58
, 27
Ni58
6
C14
7
N14
e – 6 7
p+ 6 7
n 8 7
*Different no of e–
, P+
, n
*n + p = 14, 14
Same nucleons
*Nuclear charge different
(3) Neutron
Racm your Brain
Why isobars have same physical
properties?
Ǫ
The nucleus of tritium contains
[6IPMT]
(1) 1 proton + 1 neutron

15.
Atomic
Structure
(3) Isotones
The atoms of different element which
have
the same number of neutrons.
Ex: 40
20 Ca
39
19K
n = 39 –
19
n = 20
n = 40 –
20
n = 20
(4) Isodiaphers
The atoms of different element which have
the same difference of the number of
Neutrons and protons.
5
B11
6
C12
7
N15
8
O17
p+ 5 6 7 8
n 6 7 8 9
e– 5 6 7 8
(n
–
p+
)
1 1 1 1
Ǫ
An isotone of 32
Ge76
is
[6IPMT]
(1) 32Ge77 (2) 33As77
(3) 34
Se77
(4) 36
Sc77
Ǫ
The number of protons,
neutrons 71
and electrons in 175
Lu , respectively
are :
[NEET-2020]
(1) 71, 104 and 71
(2) 104, 71 and 71
(3) 71, 71 and 107
(4) 175, 104 and 71
Racm your Brain
What is the difference
between
isotopes and isobiaphers?

16.
Atomic
Structure
(5) Isosters
The molecules which have
the
number of atoms and electrons.
s a m e
CO2
N2
O
Total
Atoms
3 3
No. of e–
6 + 8 × 2
= 22 e–
7 × 2 + 8
= 22 e–
(6) Isoelectronic
Atoms / ions/ molecules having similar no.
of e –
Ex:
H2
O and NH3
10e–
10e–
Cl–
and Ar
18e–
18e–
B3
N3
H6
and C6
H6
42e–
42e–
Ǫ
Be2+
is isoelectronic with
[6IPMT-2014]
(1) Mg2+
(3) Li+
(2) Na+
(4) H+

17
.
Atomic
Structure
(7) Relative 6bundance
Isotopes of an element occur in
different percentages in nature, which is
termed as relative abundance.
Using this relative abundance the
average atomic mass of the element can be
calcualted.
Ex: 35
Cl and 37
Cl
(75% and 25% abundancve respectively)
W6VE
⚫ A wave motion is a mean of transfer of
energy from one point to another point
without any conveying of matter between
the points.
⚫ When we throw the piece of stone
particles
on water surface in a pond, we
observe
circles of ever increasing radius, till
these strike on the wall of the pond.
⚫ When we put piece of cork on the surface
of this water, we observe that the cork
moves up and down as the wave passes,
but the piece does not travel along with
the waves.
6 wave is characterized by six characterstics
(1) Wavelength ()
⚫ It is defined as distance between two
nearest crest or nearest through. It is
measured in term of a Å (Angstrom), pm
(Picometer), nm (nanometer), cm
(centimeter), m (meter)
1Å  1010
m, 1Pm  1012
m
1nm  109
m, 1cm  102
m
(2) Frequency ()
⚫ Frequency of a wave is a number of
waves which pass through a point in 1
sec. it is measured in term of Hertz (Hz),
sec–1
, or cycle per second (cps)
1 Heartz = 1 sec–1
= 1 cps.
Concept Ladder
The upper most point
of the wave is called
crest and the lower
most point is called
trough.
Racm your Brain
What frequency is harmful
to humans?
Ǫ
The frequency of an
electromagnetic radiation is 2 ×
106
Hz. What is the wavelength in
metres
[6IPMT]
(1) 6.0 × 1014
(3) 1.5 × 102
(2) 1.5 × 104
(4) 0.66 × 10–2

18.
Atomic
Structure
(3) Time period (T)
⚫ Time taken for one complete oscillation
of wave is known as period (T). Time
taken by wave to travel a distance
equal to one wavelength. If C is the
speed of wave, then
C 

T
(4) Wave Number (  )
⚫ Number of wavelength per unit
length.
 
1

(5) 6mplitude (6)
⚫ It is the height of depth or crest of a
through of a wave.
(6) Velocity (C)
⚫ It is defined as distance covered by a
wave
in 1 sec.
 
C 
Electromagnetic Waves (EMW)
⚫ It contains electric and magnetic field.
⚫ Energy is always transferred in the form
of waves with the speed of light (3 × 108
m/s).
⚫ It’s a pure energy waves.
⚫ It does not contain mass no medium
is required for transmission.
⚫ Direction of propagation is
perpendicular
from both electric field and magnetic field.
⚫ There are various types of
electromagnetic waves (radiation) which
differs from one another in wavelengths.
Ex: Cosmic Rays, g-rays , X-rays, U.V, visible,
I.R, Micro, Radio.
Racm your Brain
What is the speed of EMW
through the vacuum?
Concept Ladder
The amplitude of a
wave is related to the
amount of energy it
carries. The
sound is
louder if
increases,
perceived as
the
amplitude and
softer if
the amplitude
decreases.
Concept Ladder
Near infrared waves
are used in remote
controls. Far infrared
waves are radiant heat.
Gamma rays have the
greatest energy.
Racm your Brain
What would happen if there
was no electromagnetic
spectrum?

19.
Atomic
Structure
Electromagnetic Spectrum
⚫ Arrangement of varous types of
electrognetic radiations in order of their
increasing (or decreasing) wavelengths or
frequencies is known as electromagnetic
spectrum.
Maxwell Theory of Electromagnetic Wave
⚫ All the radiations have wave nature
which explains interference (linear
superposition) and diffraction.
⚫ They consist of oscillating electric
and magnetic field perpendicular to each
other and to the direction of propagation.
⚫ All the radiations (radio waves, micro
waves, infra red waves, visible, UV, X-rays,
g-rays) travel at the speed of light in
vaccum.
⚫ Energy of electromagnetic wave is
proportional to amplitude and not linked
with frequence of waves.
Racm your Brain
How important are EM waves
in our lifes?
with
Ǫ
Electromagnetic radiation
maximum wavelength is
[6IPMT]
(1) Ultraviolet
(2) Radiowave
(3) X-ray
(4) Infrared

20.
Atomic
Structure
Limitations of Maxwell Theory of Electromagnetic
Wave
⚫ Phenomenon of black body
radiations.
⚫ Photoelectric effect.
⚫ Line spectra of atoms
SPECTRUM
⚫ When light coming from a source is
scatterd by a prism, light of different
wavelength are deviated through different
angles and get separated. This
phenomenon is known as dispersion and
such a dispersed light may be received
on photographic plate or it’s may be
viewed directly by eye. A collection of
dispersed light giving its wavelength
composition is called a spectrum.
theory
not
Racm your Brain
Why was
Maxwell accepted?
Concept Ladder
Most of the light in
the universe is invisible to
our
eyes. The light we can see,
made up of the
individual colors of the
rainbow.

21.
Atomic
Structure
Types of Spectrum
(1) Emission spectrum
Any substance on heating gets excited
on absorbing energy or at a very high
temperature or afterwhich radiations are
emitted from the
substance. The radiations when
analysed with the help of spectroscope,
spectral lines are obtained.
Emission spectrum may be classified as :
Concept Ladder
Emission
is
referred
emission
spectroscop
y to as
optical
spectroscopy
because of the light
nature of what is being
emitted.

22.
Atomic
Structure
(i) Continuous spectrum
When sunlight is passed through a prism,
it gets dispersed into continuous bands of
different colours.
(ii) Line spectrum
If the radiations obtained by the excitation
of a substance are analysed with the help
of a spectroscope, a series of thin bright
lines of specific colours are obtained.
There is dark space in between two
consecutive lines. This type of spectrum is
called line spectrum or atomic spectrum.
(2) 6bsorption spectrum
When white light of an incandescent
substance is passed through any
other substance, this substance absorbs
the radiations of particular wavelength
from the white light. On analysing the
transmiited light we obtain a spectrum in
which dark lines of specific wave lengths
are observed. These lines constitute the
absorption spectrum.
a
continuous
Racm your Brain
Why is the sun
spectrum?

23.
Atomic
Structure
PL6NCK'S U6NTUM
Ǫ THEORY
Diffraction and interference are
explained
by wave nature of electromagnetic
radiation.
However, some of the observations are
given which could not be explained with the
help of even the electromagnetic theory
19th century physics (konwn as classical
physics) :
⚫ Nature of emission of radiation from
hot bodies (black -body radiation).
⚫ Ejection of electrons from metal surface
when
radiation strikes it (photoelectric effect)
⚫ Variation of heat capacities of solids is
a function of temperature.
⚫ Line spectra of atoms with special
reference to hydrogen.
⚫ According to this theory, atoms or
molecules can emit or absorb energy
only in discrete quantities (small packets)
and not in any arbitrary amount. Planck
gave name quantum to the smallest
quantity of energy that can be emitted in
the form of E.N. radiation.
Ǫuanta
Ǫuantu
m
Definition
The radiant energy
emitted
or absorbed
discontinuousl
y small
discrete
by a body
in the form
of packets.
These
packets are called
quantum.
Concept Ladder
Body can emit or
absorb energy as h,
2h ...... but it can not
emit or absorb energy in
fractional values of h
such as 1.5 h, 2.5 h.
Ǫ
Calculate the energy in
joule corresponding to light
ofwavelength 45 nm.
[NEET-2014]
(1) 6.67 × 1015
(2) 6.67 × 1011
(3) 4.42 × 10–15
(4) 4.42 × 10–18

24.
Atomic
Structure
⚫ Energy of photon is proportional to
frequency and is given by
E  
E = h; h = Planck's
constant
= 6.626 × 10–34
J sec

E =
hc
; 




 
c

⚫ A body can emit or absorb energy only
in terms of the integral multiples of
quantum, i.e.
E = n. h, where n = 1, 2, 3, .........
Blacm Body Radiation
⚫ Black body radiation phenomenon first
given
by Max Planck in 1900.
⚫ The ideal body, which emits and
absorbs radiations of all frequencies, is
called black body and the radiation
emitted by that body is known as black
body radiation.
Ǫ
The value of Planck's constant
is
6.63 × 10–34
Js. The speed of
light is 3 × 1017
nm s–1
. Which of
the given values is closest to the
wavelength
in nanometer of a quantum
of
light with frequency of 6 × 1015
s–1
?
[NEET-2013]
(1) 50
(3) 10
(2) 75
(4) 25
Concept Ladder
Blackbody is a surface
that absorbs all radiant
energy falling on it. The
term arises because
incident visible light will be
absorbed rather than
reflected, and therefore
the surface will appear
black.

25.
Atomic
Structure
3 × 108
photons of a certain light radiation are found to produce 1.5 J of
energy. Calculate the wave length of light radiation ?
ETotal
= n × {h}
ETotal





= n ×
hc

 
8 6.6  10–34
 3  108

1.5 J = 3  10 

 = 3.96 × 10–
17
m
Ǫ
1
6
1
100 watt bulb emits monochromatic light of wave length = 400 nm.
Calculate the no. of photons emitted per second?
(1) 5 × 1020
(2) 3 × 1020
(3) 4 × 1020
(4) 2 × 1020
(4)
t = 1
s
Etotal =
nhc

n  6.6  10–34
 3  108
100 =
400  10–9
n = 2 × 1020
no. of photons per
second
2
Ǫ
62

26.
Atomic
Structure
PHOTOELECTRIC EFFECT
The ejection of electrons when light
of certain minimum frequency called as
threshold frequency is incident on a metal
surface is called as photoelectric effect.
Threshold Frequency
The minimum frequency of incident
light which can cause photo electric emission
i.e. this frequency is just able to eject
electrons with out giving them additional
energy.
Worm Function
The minimum quantity of energy which is
required to remove an electron to infinity
from the surface of a given solid, usually a
metal.
Incident energy = Work Function () +
K.E.max
Ei
=  + (K.E.)max
2
0 e
1
h = h  m
 2
Where me is mass of the electron and v is
the velocity associated with the ejected
electron.
Some facts of Photoelectric Effect
⚫ There is zero time lag between incidence
of light and emission of photoelectrons.
Concept Ladder
The maximum
kinetic
energy of
photoelectrons depends on
the frequency of incident
radiation; but, it is
independent of the
intensity of light used.
best
for
Racm your Brain
Which metal is
photoelectric
effect?
Ǫ
In photoelectric effect, the
kinetic energy of photoelectrons
increases linearly with the
[6IPMT]
(1) Wavelength of incident light
(2) Frequency of incident light
(3) Velocity of incident light
(4) Atomic mass of an element

27.
Atomic
Structure
⚫ For the emission of photoelectrons,
frequency of incident light must be equal
to or greater than the threshold frequency.
⚫ Rate of emission of photoelectrons from
a surface of metal is directly proportional
to the intensity of incident light.
Concept Ladder
The minimum potential
at
which the
photoelectric current
becomes zero is called
stopping potential.
The threshold frequency 0
for a metal is 6 × 1014
s–1
. Calculate the
kinetic energy of an electron emitted when radiation of frequency  = 1.1
× 1015
s–1
hits the metal.
1
K.E. = me
V2
= h ( –
0
)
2
 K.E. = (6.626 × 10–34
) (1.1 × 1015
– 6 × 1014
)
 K.E. = (6.626 × 10–34
) (5 × 1014
) = 3.313 × 10–19
J
3
Ǫ
63
BOHR'S 6TOMIC MODEL
It is a quantum mechanicla model. This
model based onquantum theory of radiation
and classical law of physics. This model
explain the stability of the atom and
emission of sharp spectral lines.
Racm your Brain
Do emitted photoelectrons
have same kinetic energy?

28.
Atomic
Structure
Bohr Model's Postulates
⚫ Atom has a central core nucleus where
the protons and neutrons are present. Size
of the nucleus is very small.
⚫ Negatively charged electron are
revolving around the nucleus in the same
wasy as the planets are revolving around
the sun. The path of electron is circular.
⚫ Electrons can revolve only in those
orbits whose angular momentum (mvr) is
integral multiple of h .
2
i.e.
mvr 
nh
2

⚫ Absorption or emission of radiation by
an atom takes place when an electron
jumps from on stationary orbit to another.
⚫ The radiation is emitted or absorbed as a
single quantum (photon) whose energy is
equal to the difference in energy of the
electron in the two orbitals involved.
Thus, e = h, where h = Planck's
constant and n = frequency of the
radiant energy. Hence the spectrum of the
atom will have certain fixed frequence.
⚫ The lowest energy state (n = 1) is called
the group state. After absorption of
energy,electron gets excited and jumps to
an outer orbit. It has to fall back to a
lower orbit with the release of energy.
Racm your Brain
Angular momentam is
integral multiple of h/2. Why not
fractional multiple is possible?
Concept Ladder
Bohr's theory
satisfactorily explains the
spectra of species having
opne electron, viz. H, He+
, Li2+
etc.
Concept Ladder
mvr 
nh
2

n2
r  0.529 
Z
A
Z
v  2.188  106

n
m / sec
n2
E  13.6 
Z2
eV / atom

29.
Atomic
Structure
Radius of the Bohr's Orbit
Let us consider an electron of mass 'm'
and charge 'e' is revolving around nucleus
having charge 'Ze' (where Z is atomic
number & e is charge) with a linear or
tangential velocity of 'v'. Further, let us
consider that 'r' is radius of orbit in which
electron is revolving.
According to Coulomb's
law,electrostatic force of attraction (F)
between moving electron and nucleus is –
F 
KZe2
r2
Where : K = constant 
1
 9  109
Nm2
/
C2
r
4o
and the cnetripetal force F

mv2
Hence mv2

KZe2
r
mr
or
v2

KZe2
mr
......(1)
From the postulate of
Bohr,
mvr 
nh
or
v2

2
n2
h2
42
m2
r2
......(2)
mvr=
1h
2

2h 3h
4h
5h 2
2
2
2
Ǫ
If r is the radius of the first
orbit, the radius of nt h
orbit of
H-atom is given by
[6IPMT]
(1) rn2
(3) r/n
(2) rn
(4) r2
n2

30.
Atomic
Structure
From equation (1) and
(2)
 r

n2
h2
42
mKZe2
On putting value or e, h,
m,
Z
r  0.529 
n2
A
Velocity of an electron in Bohr's Orbit
The total energy of an electron is revolving
in a particular orbit is —
mvr 
nh v 
nh
2
2mr
Purring the value of r in above equation
then
v 
nh 
42
mZe2
2mn2
h
2
v 
2Ze2
nh
on purring the values of e
and h
v  2.188  106

Z
m / sec
n
Calculation of energy of an electron
The total energy of an electron revolving in
a particular orbit is —
T.E. = K.E. + P.E.
1
r
The K.E. of an electron 
2
mv2
and the P.E. of an electron  
KZe2
Hence, T.E.  2 r
mv2

1 KZe2
......(3)
or
r r
r2
But mv2

KZe2
mv2

KZe2
Substituting value of mv2 in the equation
(3)
T.E.  
KZe2

KZe2
KZe2
r
r
2r
KZe2
So, T.E.  
2r
Ǫ
The energy of second Bohr orit
of the hydrogen atom is –328 kJ
mol– 1
; hence the energy of
fourth Bohr orbit would be
[6IPMT]
(1) – 41 kJ mol–1
(2) – 82 kJ mol–1
(3) – 164 kJ mol–1
(4) – 1312 kJ mol–1
Racm your Brain
What will be the value of
Kinetic energy and Potential
Energy at n = ?
∞
Concept Ladder
Bohr's atomic
model explained the
stability of an atom.
According to Bohr, an
electron revolving in a
particular orbit cannot
lose energy. Therefore,
emission of radiation is not
possible as long as the
electron remains in one of
its energy levels and hence
there is no cause of
insatbility in his model.

31.
Atomic
Structure
Substituting value of 'r' in the equation of
T.E.
E  
kZe2

42
Ze2
mk
 
22
Z2
e4
mk2
2 n2
h2
n2
h2
Thus, the total energy of an electorn in nt h
orbit is given by
E  2 2 2 4 2
n2
H2
n2
n2
 Z e
mk
 13.6 
Z2
eV / atom
 21.8  1019

Z2
J / atom
n2
 313.6 
Z2
Kcal / mole
Relationship between P.E., K.E. & T.E.
T.E. 
P
.E.
 K.E.
P.E.  
KZe2
, K .E . 
1 KZe2
, T.E.  
1 KZe2
r 2 r
2 r
2
Ǫ
Based on equation E = – 2.178
×
10–18 J(Z2
/n2
), certain conclusions
are written. Which of them is
not correct?
[NEET-2013]
(1)) Equation can be used
to calculate
change in
energy
when electron changes
orbit.
(2) For n = 1, electron has a
more negative energy
than it does for
n = 6 which means that
electron is more loosely
bound
(3) Negative sign in equation
simply means that
energy of electron
bound tot he nucleus is
lower than it would
be if electrons
were at the infinite
distance from nucleus.
(4) Larger the value of n,
larger would be the orbit
radius..

32.
Atomic
Structure
Calculation of the number of revolutions of the
electron in an orbit per second
Number of revolutions per sec. 
velocity of the electron

v

nh

1
2r 2mr
2r

nh
2

Circumference of the orbit
[On substituting the value of v from mvr 
nh
]
42
mr2
No. of revolutions per second


2 2 2
n2
h2
4 mr 4
m





nh nh  42
mze2
k

2

42
mz2
e4
k3
n3
h3
Ionization Energy / Ionization potential
Minimum energy required by an electron to
leave the ground state:
For Ex: H
(IP)H
= – (13.6) = 13.6
Separation Energy
Minimum amount of energy required to
escape an e–
from an excited state
2 


 
3 


 
Ǫ
In hydrogen atom, energy of
first excited state is –3.4eV.
Then find out K.E. of same orbit
of hydrogne atom.
[6IPMT]
(1) +3.4 eV
(3) –13.6 eV
(2) +6.8 eV
(4) +13.6 eV

33.
Atomic
Structure
Failure of Bohr Model
⚫ The theory was very successful in
predicting and accounting the energies of
line spectra of hydrogen i.e. one electron
system. It could not explain theline spectra
of atoms containing more than one
electron.
⚫ Theory does not explain the presence
of multiple spectral lines.
⚫ Theory does not explain splitting of
spectral lines in magnetic field (Zeeman
effect) and in electric field (Stark effect).
Intensity of spectral lines was also not
explained by the Bohr atomic model.
⚫ This theory could not explain
uncertainty principle.
The Bohr-Sommerfeld Theory
For explaining fine structures of
spectral lines Sommerfeld introduced two
modification in Boht's theory.
⚫ According to Sommerfeld, path of
an
electron around nucleus, is an ellipse
with nucleus at one of the foci. Circular
orbit is special case of the ellipse.
⚫ Velocity of electron moving in an
elliptical orbit varies at different parts of
the orbit. This causes relavitistic variation in
mass of moving electron. Therefore, he
took into account relativistic variation of
mass of electron with velocity. Therefore,
this is known as relativistic atom model.
Limitations of the Bohr-Sommerfeld Theory
⚫ Sommerfeld's theory was able to give
an explanaiton of the fine structure of
the spectral line of hydrogen atom. But
he could not predict the correct of
spectral lines.
Ǫ
theory
by
orbits
for
Who modified
Bohr's introducing
elliptical
electron path? [6IPMT]
(1) Rutherford
(2) Thomson
(3) Hund
(4) Sommerfel
d
Concept Ladder
Bohr-Sommerfeld
theory described the atom
in terms of two quantum
numbers, while Bohr ahd
originally used only one
quantum number.
Racm your Brain
What is the success of
the Sommerfeld model?

34.
Atomic
Structure
Which one is in correct about the angular
momentum?
(1) mvr 
h
(2) mvr 
h
(3) mvr 
3h
2 
2
(4) mvr 
3h
4

(4)
(1) n =
1
(2) n = 2
2h
2

(3) mvr =
3h
2

4

n = 3 (4) mvr =
3h
n =
3/2
4
Ǫ
64
For H-atom, calculate
radius
Z =
1
  



n H
n2

r  0.529 Å
Z


1 H
r   0.529 12
 Å


2 H
r   0.529 22
 Å


3 H
r   0.529 32

Å


n H
r   0.529 n2
 Å


2
n 1
r  n
r
(r3
)H
= (9r1
)H
rn
=
16r1
r5
= 25r1
5
Ǫ
65
()
PE = – x = 3.02
KE =
x
2
= 3.02
 .51
2
The P.E. of an e–
in the H–
atom is –3.02 eV. Then find
out
(1) K.E. (2) T.E. (3) Orbit
(4) Radius
6
Ǫ
66

35.
Atomic
Structure
T.E = –
x

–3.02
 –1.51
2 2

–13.6Z2
–
1.51
n2
–13.6
 –1.51
n2
n2
= 9 ; n = 3
3
1
[r = 0.529 ×
9
Å ]
If I.P of H atom is ‘x’ eV. Then calculate the required energy to exited e–
from 2nd
to 3rd
orbit :
(1)
2
x (2)
1
x (3)
5
x
36 8
36
(4)
7x
3.6
7
Ǫ
(3)
I.P = – E1
= x
E1
= – x = –13.6
x = 13.6
      

3 2

E –
E
  –13.6  –  –13.6   13.6  1

1 
 x
 5 
32  
22  
4 9
 
36

67
HYDROGEN SPECTRUM
⚫ Hydrogen spectrum is an example of
atomic or line emission spectrum.
⚫ Whenever an electric discharge is passed
to hydrogen gas at low pressure, a bluem
light is emitted. The light shows
discontinuous line spectrum of several
isolated sharp lines through prism.
⚫ All these lines of H-spectrum have
Lyman,
Balmer, Paschen, Brackett, Pfund and
Humphrey series.
⚫ Wavelength of various H-lines
Rydberg introduced the following
expression,





1 
1 2

2 2
 c
 
1


 R
 1
n n
R is a Rydberg's constant its value is
109,67800 m–1.
Concept Ladder
The spectral series are
important in astronomy
for detecting the
presence of hydrogen and
calculating red shifts.

36.
Atomic
Structure
Concept Ladder
Maximum number of
lines produced when an
electron
jumps from nth level to
group
level 
nn  1
2
Racm your Brain
Why the line spectra of
two elements are not
identical?
Total Number of spectral Lines
= n2 – n1  n2 – n1  1 2
Ex In the H- atom if an e–
moves from 6t h
to
3rd
orbit by transition is multi steps. Then
find
(i) Total no of spectra lines
(ii) No of lines in visible region
(iii) No of lines in UV region
(iv) No of lines is Brackett series
(v) No of lines in Infrared series
(vi) No of lines in Paschen series
Sol Lyman  O,
(IR) Paschen  3
P– Fund  1 (IR)
Balmer  O (UV)
Brackett  2 (IR)

37.
Atomic
Structure
SERIES 7 6 5 4 3 2
Lyman 6 5 4 3 2 I (UV)
Balmer 5 4 3 2 1 0 visible
Paschen 4 3 2 1 0 Infared Ray
(1)
Brackett 3 2 1 0
P-fund 2 1 0
Humphry 1 0
(21) (15) (10) (6) (3)
A certain electron transition from an excited state of H-atom in one or
more steps give rise of 5 lines in U.V region. Then how many lines does this
transition produce in IR region?
(1) 6 (2) 7 (3) 8 (4) 9
(1)
5  U.V
2  IR
IR =  4 UV
3 visible
4  visible 3  IR
1  IR
2 IR 1 IR
8
Ǫ
68
The e–
in a hydrogen atom transition from the Bohr orbit 5 to the
orbit 2. Calculate the wavelength of photon emitted during transition.
Given h =
6.6×10–34
J/sec.
C= 3× 108
m/s. RH
= 2.18 × 10–19
J
R 
RH
hc



2 2
 1 2

1
 RZ2  1
–
1

n n
9
Ǫ
69

39.
Atomic
Structure
Ex Emission Line Any particular series
spectrum Lyman series

hc
E  h

Lyman last line (∞1)
E  max  
max
*


 2 2
 1 2

1
 Rz2  1
–
1

n n
n  n
( ) –1
2 1
R = 109678 cm
= 1097.0 cm–1
1
 912Å
R


 

2 2
1
 RZ2  1
–
1

2 3
10
Ǫ Calculate the  of 1st
lined and last line of Balmer series in H-
spectrum ?
610
 
= R 
 1
–
1


4
9

1

5R

36
 


36

 1  
36

912Å
5 R

5
Last line
  L


  
2 2
1
 R
 1
–
1

2 
 

1
 R 1


4

 
4 R
= 4 × 912
Å
Ǫ
Which of the following series
of transitions in the spectrum
of hydrogen atoms falls in visible
region:
[NEET-2019]
(1) Brackett series
(2) Lyman series
(3) Balmer series
(4) Paschen series

40.
Atomic
Structure
W6VE MECH6NIC6L MODEL OF 6N 6TOM
The model consists of following :
(A) de-Broglie concept (Dual nature of
Matter)
(B) Heisenberg’s Uncertainty principle.
DU6L N6TURE OF M6TTER (W6VE N6TURE OF
ELECTRON)
 
1
p p
or  =
h
(Here h = Planck’s
constant,
mv
  =
h
We know that according to
Bohr
theory, mvr =
nh
2

or 2pr
=
nh
mv
 2pr = n
nh
mvr =
2
p = momentum of electron)
 Momentum (p) = Mass (m)×Velocity (v)
Ǫ
In hydrogen atom, the de
Broglie wavelength of an electron
in the second Bohr orbit is
[Given that Bohr radius, a = 52.9
pm]
(1) 211.6 pm
(3) 52.9 
pm
0
[NEET-2019]
(2) 211.6  pm
(4) 105.8  pm
Concept Ladder
Circumference of orbit
is equal to integral
multiple of wavelength (),
i.e., 2r = n.

41.
Atomic
Structure
Heisenberg uncertainty principle
When an electron is considered to be a
wave as suggested by de-Broglie, it is not
possible to identify the exact position and
velocity of the electron more precisely at a
given instant since the wave extends
throughout a region of space.
Dx.Dp 
h
or Dx. mDv 
h
or Dx.Dv 
h
4 4 4m
t
4
or Dt×Dx×
p

h
Where h is Planck’s
constant.
F × Dt × Dx 
h
4

E×
t 
h
4

From above description it is clear that
according to de-Broglie there is similarity
between wave theory and Bohr theory.
A ball weight 25 g moves with a velocity of 6.6×104
cm/s then find out the
de-
Broglie  associated with it.

=
= 4
h 6.6  1034
 107
erg
mv 25  6.6  10 cm /
s
s = 0.04×10–31
cm = 4 × 10–33
cm
11
Ǫ
611
12
Ǫ If the uncertainity in position of a moving particle is 0 then find out

p.
Ǫ
Which one is the wrong
statement?
[NEET-2017]
(1) The unceratnity principle is
E  t 
h 4

(2)Half filled and fully filled
orbitals have greater stability due
to greater exchange energy,
greater symmetry and more
balanced arrangement.
(3)Then energy of 2s-orvbital is
less than the energy of 2p-orbital
in case of hydrogen like atoms.
(4) de-Broglie's wavelength is
given
 
h
mv
by where m = mass
of
particle, v = group velocity of the
a particle
Racm your Brain
Why possiblity of the electron
in the nucleus is zero?

42.
Atomic
Structure
W6VE MECH6NIC6L MODEL OF 6TOM
⚫ This model was developed by
Erwin
Schrodinger in 1926.
⚫ This atomic model is based on
particle
and wave nature of electron is known as
wave mechanical model of the atom.
⚫ This model describes the electron as a
three- dimensional wave in the electronic
field of positively charged nucleus.
⚫ Schrodinger derived an equation
which described wave motion of an
electron. The differential equation is :
dx2
dy2
dz2
h2
d2


d2


d2


82
m
E  V   0
where x, y and z are cartesian co-ordinates
of the electron
m = mass of the elecron
E = total energy of electron
V = potential energy of
electro n h = planck's
constna
 = wave function of the
electron
2
    0
82
mK.E.
h2
Significance of 
Wave function is regarded as the
amplitude function expressed in terms of
coordination x, y and z. Wave function can
have negative or
h
D x Dp 
4
or Dp 
h
4x
h
or Dp
 4
0
or Dp  
612
Ǫ
The uncertainty in momentum
of an electron is 1 × 10–5
kg m/s.
The uncertainty in its position will
be
[6IPMT]
(1) 5.27 × 10–30
m
(2) 1.05 × 10–26
m
(3) 1.05 × 10–28
m
(4) 5.25 × 10–28
m
Concept Ladder
Berhaviour of electrons
and other microscopic
particles, a new branch
of science called quantum
mechanics was developed.

43.
Atomic
Structure
positive values depending upon the values
of coordinates.
Significance of 2
2
is a probability factor. It describes
the probability of fi8nding an electron
within a small space. The spacein which
there is maximum probability of finding an
electron is termed as orbital.
Variations of Radial Wave Function (R)
(i) Plots of radial wave function R against the
distance r
The variation of the radial part of the
orbital wave funcitons for 1s, 2s and 2p
orbitals. The radial funcitn value changes
form positive to zero then to negative.
The region where this function reduces to
zero is called nodal surfaces or simply
nodes.
(ii) Radial probability density (R2
)
Square of radial wave function, R2
, is
the measure of the probability of
finding the electron in a unit volume
around a particular point and is called
probability density.
Ǫ
The graph between ||2
and r
(radial distance) is shown
below. This represents
(1) 3s orbital
(3) 2p
orbital
[6IPMT]
(2) 1s orbital
(4) 2s orbital
Racm your Brain
What is the significance of
the
wave function ?

44.
Atomic
Structure
s- Subshell
(iii)Radial distribution function, (4r2
R2
)
Probability density of finding electron at
a point at a distance r form the nucleus.
Since the atoms have spherical symmetry, it
is more useful to discuss the probability
of finding the electron in a sphericla shell
between the spheres of radii r + dr and r.
Node
Nodal point
(=0) Nodal
plane Nodal
surface
Nodal plane / Angular Node
= l Nodalsurface / Radial
Nodes = n – l –
1
Radial Nodes
Angular
Nodes
Radial Nodes 
Racm your Brain
How do you find number of
nodal
planes in a molecular orbital?
Ǫ
Orbital having 3 angualr nodes and
3 total nodes is
[NEET-2019]
(1) 5p
(3) 4f
(2) 3d
(4) 6d
Definition
The plane / surface in
which probability of finding
electron is zero is called as
Node.

45.
Atomic
Structure
Radial Nodes (n –l –1)
The number of nodes
is always one less than
the principal quantum
number
: Nodes = n - 1. In the
first electron shell, n = 1.
the 1s orbital ha no nodes.
The 3s, 3p and 3d orbitals
have two nodes, etc.
Angular Node l = 0
RN = n – l – 1 = n =
2
C
o
n
c
e
p
t
L
a
d
d
e
r
Definition
To obtain the information about
an electron Identification
numbers are required. These
numbers are called Ǫuantum
numbers.
n – 2 [RN]
1s 0 2px
0
2s 1 3px 1
3s 2 4px 2
U6NTUM
Ǫ NUMBERS
⚫ Ǫuantum numbers are to specify and
display to complete information about
size, shape and orientation of the orbital.

46.
Atomic
Structure
Types of uantum
Ǫ Numbers
(1) Principal Ǫuantum Number
(n)
(2) Azimuthal Ǫuantum Number
(l)
(3) Magnetic Ǫuantum Number (m)
(4) Spin Ǫuantum Number (s)
(1)
⚫
Principal uantum
Ǫ Number (n)
It is the most important quantum
number
as
it determines the size and to large extent
the energy of the orbital.
⚫ The average energy of the electron is
directly proportional to the principal
quantum number.
⚫ The size of an orbital will increase an
increase in the principal quantum number.
⚫ Maximum number of electrons in a shell
is given by 2n2
.
Exclusive identity of an e–
*n  name, size energy of
shell
n = 1
n = 2
n = 3
n = 4
n = 5
n = 6
K Shell
L
M
N
O
P
6ngular momentum (J/L) mvr=
nh
2

J > J > J
3 2
1
Max. No of e–
in a shell =
2n2
K n = 1 (2)
L n = 2 (8)
M n = 3 (18)
N n = 4 (32)
Ǫ
The orientation of an atomic
orbital is governed by
[NEET]
(1) principal quantum number
(2) azimuthal quantum number
(3) spin quantum number
(4) magnetic quantum number
Concept Ladder
Ǫuantum number are
important because they
can
be
determine
configuration,
of
the elctron
probale
electrons,
energy,
atomic
location
ionization
radius.

47.
Atomic
Structure
Value of l 0 1 2 3
Sub-shell notation s p d f
Number (2l + 1) of orbitals 1 3 5 7
(2) 6zimuthal uantum
Ǫ Number (l)
⚫ This is also known as orbital
angular momentum or subsidiary quantum
number. (l)
gives the in
which the
⚫ Azimuthal quantum
number information about
subshell electron is located.
⚫ Orbital Angular Momentum


 
2

l l  1 h 
l l  1 k
s l = 0
→
p l = 1
→
d l = 2
→
f l = 3
→
0

k


k


k

2
6
12
Concept Ladder
The
name
'azimutha
l
quantum number' for l
was originally
introduced by
Sommerfeld, who
refiend Bohr's semi-classical
model by replacing
circular orbits with
elliptic ones. The spherical
orbitals were similar (in
the lowest-energy state) to
a rope oscillating in a large
'horizontal' circle.

48.
Atomic
Structure
⚫ Max. No. of e–
in a subshell = (4l +
2)
s l = 0 2 for any given n, l
will
be
p l = 1 6
from [0 to n =–1]
d l = 2 10
f l = 3 14
Shape of subshell
⚫ s, p, d and f are taken from spectroscopic
terms sharp, principal, diffuse, and
fundamental, respectively.
⚫ In a multi electron atom, the energy
associated with an electron depends both
on n and l.
n l Subshell notation
1 0 1s
2 0 2s
2 1 2p
3 0 3s
3 1 3p
3 2 3d
4 0 4s
4 1 4p
4 2 4d
4 3 4f
Concept Ladder
The magnetic
quantum number describes
the energy
subshell and yields
levels available within a
the
projection of the
orbitgal angular
momentum along a
specified axis.
Ǫ
The following quantum number
are possible for how many orbitals?
n = 3, l = 2 and m = +2
[NEET]
(1) 1
(3) 3
(2) 2
(4) 4
Ǫ
What is the maximum numbers
of electrons that can be
associated with the following set
of quantum numbers?
n = 3, l = 1 and m = –1
[NEET-2013]
(1) 4
(3) 10
(2) 2
(4) 6

50.
Atomic
Structure
(3) Magnetic uantum
Ǫ Number (m)
⚫ Magnetic quantum number was proposed
by Land.
⚫ Themagnetic quantum number
gives
information about the spatial
orientation of orbitals. These different
orientations are called orbitals.
⚫ It is denoted by m and its value depend
on l values.
⚫ The possible value of m range from -l
through 0 to +l.
⚫ Total number of orbitals = (2l + 1)
⚫ Number of obitals in a shell is n2
.
(4) Spin uantum
Ǫ Number (s)
⚫ In 1925, George Uhlenbeck and
Samuel Goudsmit proposed the presence
of the fourth quantum number and
depicted it as the elctron spin quantum
number.
⚫ Tw
1
o electrons that have different s,
values

2
and both electrons have opposite
spins.
⚫ An orbital can hold maximum two
electrons.
⚫ Spin angular momentum is depicted by
the
symbol s
. The value of
s
:
s
  s s  1 h
2

Ǫ
The correct set of four
quantum number for the valence
electron of rubidium atom (Z = 37)
is
[NEET-2012]
(1) 5, 1, 1, +1/2
(3) 5, 0, 0, +1/2
(2) 6, 0, 0, +1/2
(4) 5, 1, 0, +1/2
Concept Ladder
Zeeman effect : Splitting
of lines of atomic
spectrum in magnetic field.
Stark effect : Splitting
of lines of atomic
spectrum in electric field.
Ǫ
Which of the following pairs
of d-orbitals will have electron
denstiy along the axis?
[NEET-2012]
(1) dz
2 , dxz
(2) dxz
, dyz
z2 xy
(3) d , d (4) d ,
d
x2
y2
x2
y2

51.
Atomic
Structure
⚫ If all the electrons in an atom or molecule
are paired. They behave as dimagnetic
substance. It is weakly repelled by the
magnetic field.
⚫ If atoms or molecules of a substance
have one or more unpaired electrons, it
behaves as a paramagnetic substance. it
is weakly attracted by the magnetic field.
⚫ Magnetic moment

nn  2 BM
Where, n = number of unpaired
electrons
Difference between orbit and orbital
Orbit Orbital
1. An orbit refers to
the circular path in
which an electron
revolves around the
nucleus.
1. An orbital refers
to the region of
space having the
maximum
probability of
finding an electron
around the nucleus.
2. An orbit
represents the
motion of an
electron arougn
the nucleus in a
plane.
2. An orbital
represents the
motion of an
electron around
nucleus in
three- dimensional
space.
3. An orbit (n) can
accommodate
a
maximumn of
2n2
electrons.
3. An orbital can
accommodate
a
maximum of
two electrons.
4. Orbits are
desinated as K, L, M
etc. or 1, 2, 3etc.,
from the nucleus
outwards.
4. Orbitals are designated
as dxy, dyz, dzx, dx2-y2, dz2
,
px
, py
, pz
etc.
5. Orbits are circular
in shape.
5. Orbitals have
different shapes, e.g.,
s orbitals are
spherically
symmetrical
whereas p
orbitals are
dumbbell shaped.
Concept Ladder
The spin of electrons
is responsible for most
of the magnetic
properties of
atoms, molecules, or
ions. Due to their spin,
electrons behave as tiny
magnets.

53.
Atomic
Structure
RULES FOR FILLING OF ORBIT6LS IN 6N 6TOM
(1) 6ufbau Principle
The principle states that electrons are
added progresively to the various orbitals
in the order of increasing energies. The
electrons first occupy the lowest energy
orbital avaiable to them and enter into
higher energy orbitals only after lower
energy orbitals are filled.
(n +l) Rule (For multi electron species)
The subshell with lowest (n + l) value is
filled up first, when two or more
subshell have same (n + l) value then
the subshell with lowest value of n is
filled up first.
(2) Hund’s Maximum Multiplicity Rule
According to the Hund’s rule orbital
available in the subshell are first filled
singly with parallel spin electron before
they begin to pair this means that pairing
of electrons occurs with the introduction
of second electron.
Concept Ladder
Afbau principle proposed
by Niels Bohr in the early
1920s, the principle was a
tool for obtaining a
picture of the atomic
constitution, i.e., the
arrangement of electrons
on orbits around the
nucleus.
Ǫ
In a given atom no two electrons
can have the same valus for all the
four quantum numbers. This is
called
[6IPMT]
(1) Hund's Rule
(2) Aufbau Principle
(3) Uncertainty Principle
(4) Pauli's Exclusion Principle.

54.
Atomic
Structure
(3) Pauli Exclusion Principle
According to this rule no two electron
in an atom can have same values of all
four quantum numbers.
Ex:
6
C12 
1s2
2s2
2p2
ELECTRONIC CONFIGUR6TION OF 6TOMS
In an atom, electrons are distributed
among various orbitals very much in
accordance with rules governing the filling of
different orbitals.
Ex Notation form of Na —
1s2
2s2
2p6
3s1
Orbital diagram from
of Na
Condensed form of Na — [Ne]
3s1
ST6BILITY ORDER OF COMPLETELY FILLED 6ND
H6LF FILLED SUB-SHELLS
The ground state electronic configuration
of
atom corresponds to the lowest energy state
and gives higher stability.
The electronic configuration of most of
the
atoms follows the basic rules.
Certain elements such as Cr or Cu do
not follow the rules because the two sub-
shells 4s and 3d slightly differ in energy, i.e.,
4s is slightly lower in energy than 3d orbital.
So the valence electronic configuration are
3d5
4s1
and 3d10
4s1
respectivley, and not 3d4
4s2
and 3d9
4s2
.
The extra stability of half-filled and fully
filled electronic configuration can be
explained in
Definition
Distribution of electrons
into orbitals of an atom is
called its electronic
configuration.
Concept Ladder
The Pauli exclusion
principle helps explain a
wide variety of physical
phenomena.
Electrons have to
stack
Px
Pz
Py within an atom, i.e.
have
n 1 2 2 different spins while at
teh
l 0 0 1 same electron orbital.
m 0 0 +1, 0, –1
s +½, – ½ +½, – ½ +½, +½
Ǫ
If n = 6, the correct sequence
for
filling of electrons will be
[6IPMT-2011]
(1) ns  (n 2)f  (n 1)d  np
(2) ns  (n 1)d  (n 2)f  np
(3)ns  (n 2 ) f  np  (n 1)d
(4) ns  np  (n 1)d  (n 2)f

55.
Atomic
Structure
terms of symmetry and exchange
energy.
Symmetrical distribution of electrons
The electronic configurations in which
all the orbitals of the same sub-shell are
either completely filled or half filled have
relatively more symmetrical distribution of
electrons. So their shielding of one another
is relatively small and the electrons are more
strongly attracted by the nucleus.
Ex: Cr
Exchage energy
Exchange means shifing of electrons
from one orbital to another within same
sub-shell. Energy gets released when
electrons exchange their positions and the
energy is called exchange energy.
For maximum number of exchanges,
the maximum the energy released and the
maximum the stabilisation.
Half-filled and fully-filled
degenerate
orbitals have more number of elctron
exchanges,
and consequently,
they energy of
stabilisation.
Ex Cr (For 3d4
4s2
)
have larger
exchange
Ǫ
The outer electronic configuratino
of
Gd (Z = 64) is
[NEET-2013]
(1) 4f5
5d4
6s1
(3) 4f3
5d5
6s2
(2) 4f7
5d1
6s2
(4) 4f4
5d5
6s1
Concept Ladder
Stability of
electronic
configuration depends on :
(1) Half-filled and Full-filled
(2) Symmetrical distribution
(3) Exchange Energy
Ǫ
The electronic configuratino of Cu
(Z
= 29) is
[6IPMT]
(1) 1s2
2s2
2p6
3s2
3p6
4s2
3d9
(2) 1s2
2s2
2p6
3s2
3p6
3d10
4s1
(3) 1s2
2s2
2p6
3s2
3p6
4s2
4p6
5s2
5p1
(4) 1s2
2s2
2p6
3s2
3p6
4s2
4p6
3d3

56.
Atomic
Structure
Cr (For 3d5
4s1
)

57.
Atomic
Structure
(1) 20 amu, 25%
(3) 40 amu, 25%
(2) 30 amu, 25 %
(4) 45 amu, 25 %
13
Ǫ If the mass of e–
is assumed to be doubled the mass of proton is
doubled and mass of neutron is halved, Then calculate the atomic wt. of
8
016
and % by which it is increased?
(1)
1e–
= negligible, 1P+
= 1 amu, 1n  1 amu
16
8 O
e –
 8  0 amu, p
 8  8 amu, n  8  8 amu
Total 16
amu Double
e–
= X, p+
= 16 amu, n = 4
amu Total 20 amu
% increment:
4
 100  25%
16
613
(1) 22 amu, 75%
(3) 20 amu, 50%
(2) 21 amu, 75%
(4) 15 amu, 60%
14
Ǫ If the mass of proton is halved, mass of neutron is tripled and mass of
e–
remains unchanged. Then calculate the atomic weight of 6
C12
and
the % increment:
614 (2)
6
e–
= 



Total
1
2
12 21 amu (% increment =
9
 100 =
75%)
15
Ǫ If an element has two isotopes forms 18 and 20 and the % abundance in
nature are 20% and 80% respectively. Then calculate the avg. atomic wt. of
element ?
(1) 21.6 (2) 19.6 (3) 20.6 (4) 22.6

58.
Atomic
Structure
(2)
18
A20
XA
X
615
Average atomic weight = 18 
20
 20 
80
 19.6
100
100
16
Ǫ The wave number of a beam of light is 400 cm– 1
calculate the
wavelength in
terms of nanometer. Also find its frequency.

 
–1
   400 cm
1
cm 
1
400 400  10–7
nm
616   400 cm1
1
 

c

3 
108 1
 10–2
m
400
(1) 3m, 0.22 m–1
(3) 4m, 0.33 m–1
(2) 3m, 0.33 m–1
(4) 1m, 0.30 m–1
17
Ǫ A radio station is broadcasting program at 108
frequency if the distance
b/w radio station and the receiver is 3 lakh meter then, how long would it
take the signal to reach the receiver. Find out the value of  and  ?
(2)
t 
300000 m 1
S
3  108
m / s 103

 
c

10+8
=
3  108
  
3
–1
1
1
   m
 3
617
orbit of He ?
(1) 0.16 × 10–8
cm
(3) 0.46 × 10–8
cm
(2) 0.36 × 10–8
cm
(4) 0.26 × 10–8
cm
18
Ǫ Radius of 1st
Bohr orbit of H–atom is 0.52 × 10–8
. Calculate the radius of
1st

59.
Atomic
Structure
618 (4)
n2
 

rn 
Z
2
1

2
1 H 1 1
2
2 2
r n / z
r1 He
n / z
 1 H
He
r  
r
=
12
/1
 2
2 12
/ 2
= – 8
0.52  10–8
 0.21  10 cm
2
Orbit Z
H 1 1
He+
1 2
r 
1
Z
19
Ǫ I.P of H–
atom is 13.6 eV.
Calculate the I.P of He+
& Li+2
H I.P = –E1
E1
= –13.6
He+
I.P = –E1
E1
= –13.6 ×
4
619
Li+2
I.P = –E1
= + (9 ×13.6) = –13.6× 9
I.P = {13.6} × Z2
20
Ǫ A photon of energy 12.09 eV is completely absorbed by a H–atom initially
in
ground state find out the orbit in which electron is revolving
(1) 2 (2) 3 (3) 4 (4) 5
620 (2)
–13.6 + 12.09 = –1.51 = –13.6
1
n2
n2
= 9
n = 3
21
Ǫ An e–
of an atom having z = 5 revolves around a nucleus the energy required
to excite the e–
from 3rd
to 4t h
orbit will be :
(1) 12.5 (2) 13.5 (3) 14.5 (4) 16.5

60.
Atomic
Structure
621 (4)
4 3
E –
E
 
–13.6  25  – 
–13.6 
25

 eV
 42   32



9
 
= 13.6 × 25  1
–
1  =
13.6 × 25 ×  7 
eV
16 
 16 
9 
= 16.5
Tricm

( ) × Z2
( ) Z =
5
3  4
10.2 1.89
0.660.31
0.66 × 25 =
16.5
22
Ǫ Compare the velocities of e–
in the first excited state of He+
and 2nd
excited state of Li+2
(1) 2 : 1 (2) 1 : 2 (3) 1 : 1 (4) 2 :
3
622 (3)  He+
Z 2
Li+2
3
n 2
1 =
1
3
n
[ V 
Z
]
(1) 75.5 (2) 86 (3) 86.5 (4) 87.5
23
Ǫ The relative abundance of two rubidium
isotopes
85 87
37 Rb 37 Rb
75% 25%
623 (2)
Find out average atomic
wt.
85 87
75% 25%
86

Atomic
Structure
Chapter Summary
 Atom is the fundamental unit of matter which is further indivisible i.e.
atom can neither be created nor be destroyed.
 Atomic theory of matter was first proposed by John Dalton.
 Discovery of subatomic particles namely electron and proton. Electron is
discovered through cathode ray discharge tube experiment. They consist of
negatively charged particles which are known as electrons. The
characteristics of cathode rays is independent of electrodes and nature of
gas present in cathode tube.
 Measurement of e/m for electron. The value of e/m has been found to be
1.7588
× 1011
C/Kg.
 Charge on an electron is 1.602 × 10–19
C.
 Mass of the e–
can be calculated from the value of e/m and the value of e.
 The characteristics of anode rays or canal rays is dependent on nature of
gas present in cathode tube. Charge to mass ratio of particle depends on
the gas from which they generated.
 Atoms are made of three particles electron, protons and neutrons.
 Rutherford’s Model: Atom of an element consist of a small positive
charged nucleus situated at the centre of the atom. Electrons are
distributed in different concentric circular paths around the nucleus called
orbits. Atomic radius is of the order 10–10 m while nucleus is 10–15
m.
 Atomic no. of an element = Total no. of protons present in the nucleus.
 Protons and neutrons present in nucleus collectively called nucleons.
 Mass no. of an element = No. of protons + No. of neutrons
 Isotopes: They have same atomic number but different atomic weight and
have
same chemical properties.
 Isobar: The different atoms which have same atomic masses but different
atomic
number are called as Isobar.
 Isotone: When elements have same number of electrons of neutron are called
as Isotones.
 Isoelectronic: Ion or atom or molecule or species which have the same
number of electrons are called as isoelectronic species.
 Electromagnetic Wave Radiation: The oscillating electrical/magnetic field
are electromagnetic radiations. Both electrical and magnetic field are
perpendicular to each other.
 Order of wavelength in electromagnetic spectrum
 Cosmic rays < -rays < X-rays < Ultraviolet rays < Visible < Infrared < Micro
waves
< Radio waves.
 Planm’s uantum
Ǫ Theory: Some important phenomena such as interference and
diffraction are generally explained by wave nature of electromagnetic radiation..

62.
Atomic
Structure
(i) Nature of emission of radiation from the surface of hot bodies (black - body
radiation)
(ii) Ejection of electrons from the surface of metal happens when radiation
strikes
it (photoelectric effect)
 Photoelectric Effect (P.E.E.):
The ejection of electrons when light of certain minimum frequency called as
threshold
frequency is incident on a matel surface is called as photoelectric effect.
Incident energy = Work function () +
K.K.max¬ Ei
=  + (K.E.)max
h = h
+ (1/2) me
v2
where me
is mass of electron and v is the
velocity associated with the ejected electron.
 Bohr’s 6tomic Model: Energy of an electron remains constant as long as it
stays in same orbit called stationary Orbit. A fixed amount of energy is
associated with each stationary orbit and hence it is called Energy level.
Energy of an electron:
T.E. 
P
.E.
 K.E.
2
T.E.  13.6 
Z2
eV / atom
n2
 Ground state (G.S.): In any single electron species n = 1 is called ground state.
 Excited state (E.S.): In single electron species n > 1 is called excited state.
 For the nth shell = (n - 1) the excited state
 Excitation energy: Energy required to excite an electron from its ground state
to any excited state is called excitation energy.
 Wave Mechanical Model of An Atom: WAVE MECHANICAL MODEL OFAN ATOM:
The Dual Nature of Matter (The Wave Nature of Electron)
De-Broglie Equation (Dual nature of matter and
radiation):
 
h

h
mc p
 
h

mv
h
2m(K.E.)
If a charged particle Ǫ is accelerated through potential difference V from rest
then
De-broglie wavelength is


h
2mQV

63.
Atomic
Structure
 The circumference of the nth orbit is equal to n times the wavelength of
the electron. 2rn = n.
 Heisenberg’s Uncertainity Principle: It is impossible to obtain
simultaneously both position and velocity (or momentum) of a microscopic
particle with absolute accuracy.
or x.v 
h
4
m
x.p 
h
or m
x.v 
h
4
4
 uantum
Ǫ Mechanical Model:
The Schrodinger Equation:
h2
h2
2


2


2


82
m
or 2
 
82
m
(E  V)   0
x2
y2
z2
 uantum
Ǫ Numbers:
(i) Principal quantum number (n): Number of orbitals present in nth shell = n2
.
The maximum number of electrons which can be present in a principal energy
shell is equal to 2n2
.
(ii) Azimuthal quantum number (l):
Number of orbitals in a given subshell = 2l + l
Maximum number of electrons in particular subshell = 2 × (2l + l)
Orbital angular momentum L 
h
2

ℓ(ℓ  1)  ħ ℓ(ℓ  1)
(iii)Magnetic quantum number (m): It describes the orientations of the
subshells. It can have values from – l to +l including zero, i.e., total (2l + l) values.
(iv) Spin quantum number (s): It describes the spin of the electron. It has values
+1/2 and –1/2.
(+) signifies clockwise spinning and (–) signifies anticlockwise spinning.
 Shape of The Orbitals:
Nodal plane and Nodal surface :- The space where probability of finding an e–
is
zero. Nodal plane = l ; Nodal surface = n - l – 1
 Stability of Completely Filled and Half-filled Subshells:
Symmetrical distribution of electrons: Due to small shielding, the electrons are
pulled closer to the nucleus and this decrease in energy leads to stability.
Exchange energy: Electrons with the same spin have a tendency to exchange
their positions when they are present in the degenerate orbitals of a subshell.
The energy released during this exchange is called exchange energy

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