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Detailed Revision Notes

Composition of atom
...
J
...
It is now believed that the atom consists of several sub-atomic
particles like electron, proton, neutron, positron, neutrino, meson etc
...


Electron (–1eo)
(1) It was discovered by J
...
Thomson (1897) and is negatively charged particle
...

(2) Cathode rays were discovered by William Crooke's & J
...
Thomson (1880) using a
cylindrical hard glass tube fitted with two metallic electrodes
...
This tube was known as discharge tube
...
Blue rays were emerged from the

cathode
...

Cathode rays

Gas at low
pressure

Cathode

TC Vaccum pump
Anode

High voltage
+
–

Discharge tube experiment for production of cathode
rays

(3) Properties of Cathode rays
(i) Cathode rays travel in straight line
...

(iii)Cathode rays consist of negatively charged particles known as electron
...

(vi) Cathode rays heat the object on which they fall due to transfer of kinetic energy to
the object
...


(viii) Cathode rays possess ionizing power i
...
, they ionize the gas through which they
pass
...

(x) They can penetrate through thin metallic sheets
...

(xii) The e/m (charge to mass ratio) for cathode rays was found to be the same as that
for an e

( 1
...
Thus, the cathode rays are a stream of

electrons
...
This is because the gases are poor conductor of electricity
...
Similarly,
fluorescent light tubes are also cathode rays tubes coated inside with suitable
materials which produce visible light on being hit with cathode rays
...
S
...
The
charge on each electron is 1
...

(5) Name of electron was suggested by J
...
Stoney
...
J
...

(6) Rest mass of electron is 9
...
000549 amu

1 / 1837 of the mass of hydrogen

atom
...

When u=c than mass of moving electron =
...
483

10

4


...
1 10 27 electrons =1gram
...
5483 mili gram
...
The minus sign on the electron in an orbit, represents
attraction between the positively charged nucleus and negatively charged electron
...

(13) The physical and chemical properties of an element depend upon the distribution of
electrons in outer shells
...
28

10

12

cm
...
17 10

17

g / mL
...
It is a
component particle of anode rays
...
These rays also termed as
positive or canal rays
...

(ii) Anode rays are material particles
...

(iv) Anode rays may get deflected by external magnetic field
...

(vi) The e/m ratio of these rays is smaller than that of electrons
...
It is maximum when gas present in the tube is hydrogen
...

(4) Charge on proton = 1
...
80

10

(5) Mass of proton = Mass of hydrogen atom= 1
...
s
...


1
...

(6) Molar mass of proton = mass of proton

Avogadro number

1
...


(7) Proton is ionized hydrogen atom (H ) i
...
, hydrogen atom minus electron is proton
...

(9) Mass of 1 mole of protons is

1
...


(10) Charge on 1 mole of protons is

96500 coulombs
...
58

4 3
r ) is
3

1
...


cm 3
...

(3) Neutron is slightly heavier (0
...

(4) Mass of neutron = 1
...
675
hydrogen atom
...

(6) Density = 1
...
00899 amu

mass of

gram / c
...


(7) 1 mole of neutrons is 1
...

(8) Neutron is heaviest among all the fundamental particles present in an atom
...
It decays as follows :
1
0n
neutron

1
1H
proton

0
1e
electron

0
0
anti nutrino

(10) Neutron is fundamental particle of all the atomic nucleus, except hydrogen or
protium
...
000546
9
...
00728
1
...
00899
1
...
602 × 10–19
– 4
...
76 × 108

+1
...
8 × 10–10
+1
9
...
660 10

27

Zero
6C

12

, i
...


kg
...
8029

0

e , 1e 0 ,

Nature

0

–

– 4
...
000548
6
<
0
...
00787

Discovered by

Anderson (1932)
Pauli (1933) and Fermi
(1934)
Chamberlain Sugri (1956)

Positive mu meson

+

+ 4
...
1152

and Weighland (1955)
Yukawa (1935)

Negative mu meson

–

– 4
...
1152

Anderson (1937)

Positive pi meson

+

+ 4
...
1514

Negative pi meson

–

– 4
...
1514

0

0

0
...

(1) Atomic number or Nuclear charge
(i) The number of protons present in the nucleus of the atom is called atomic number
(Z)
...

(iii) Atomic number = Number of positive charge on nucleus = Number of protons in
nucleus = Number of electrons in nutral atom
...

(2) Mass number
(i) The sum of proton and neutrons present in the nucleus is called mass number
...

(ii) Since mass of a proton or a neutron is not a whole number (on atomic weight
scale), weight is not necessarily a whole number
...
g
...


:  A part of an atom up to penultimate shell is a kernel or atomic core
...

protons

 Number of lost or gained electrons in positive or negative ion =Number of
charge on ion
...
(Z)
(ii) No
...
of neutrons

(iii) No
...
(A)
(iii) Physical properties

Examples
2
3
(i) 1 H , 1 H , 1 H
1

(ii)

16
8

(iii)

(iv) Electronic
configuration

O, 17 O, 18 O
8
8

35
17

37
Cl , 17 Cl

(v) Chemical properties
(vi) Position in the periodic
table
(i) Mass No
...
(Z)

(ii) No
...
of protons, electrons
(ii)
and neutrons

Isobars

(i)

40
18

40
Ar , 19 K , 40 Ca
20

130
52

Te , 130 Xe , 130 Ba
54
56

(iii)Electronic configuration
(iv) Chemical properties
(v) Position in the perodic
table
...
of neutrons

(i) Atomic No
...
, protons and
39
(ii) 19 K , 40 Ca
20
electrons
...

Isotopic No
...
,
electrons,
neutrons
...
, (i) 92 U 235 , 90 Th 231
protons,
(ii) 19 K 39 , 9 F 19

(ii) Physical and chemical (iii)
properties
...
of electrons

At
...
, mass No
...
of atoms

(i) N 2 and CO

(ii) No
...


Isosters

(iii) HCl and F2
(iv) CaO and MgS
(v) C6 H 6 and B3 N 3 H 6

Note :  In all the elements, tin has maximum number of stable isotopes (ten)
...

(1) Light and other forms of radiant energy propagate without any medium in the space in
the form of waves are known as electromagnetic radiations
...
e
...

rays,
rays,
cosmic rays, ordinary light rays etc
...
(ii)
These consist of electric and magnetic fields components that oscillate in directions
perpendicular to each other and perpendicular to the direction in which the wave is travelling
...
It is denoted by
(lambda) and is measured is terms of
centimeter(cm),
angstrom(Å),
Crest
Wavelength
micron( ) or nanometre (nm)
...
It is denoted by the symbol (nu) and is expressed in terms of cycles (or
waves) per second (cps) or hertz (Hz)
...
It is
denoted by the letter „c‟
...
e
...


3 10 10 cm / sec

c

Thus, a wave of higher frequency has a shorter wavelength while a wave of lower
frequency has a longer wavelength
...
e
...
It is denoted by the symbol (nu bar)
...


1
(v) Amplitude : It is defined as the height of the crest or depth of the trough of a wave
...
It determines the intensity of the radiation
...

Name

Wavelength (Å)

Frequency (Hz)

Source

Radio wave

3 10 14

Microwave

3 10 7

Infrared (IR)

6 10 6

Visible

7600

3800

3
...
9 10 14

X-Rays

150

0
...
1

0
...
01- zero

3 10 20

infinity

Outer space

Rays
Cosmic Rays

1 10 5

1 10 9

Alternating
frequency

6 10 6

1 10 9

5 10 11

Klystron tube

7600

5 10 11

3 10 7

3
...
9 10 14
2 10 16

current

of

high

lamps

with

Incandescent objects
Electric bulbs, sun rays
Sun rays, arc
mercury vapours

Atomic spectrum - Hydrogen spectrum
...
It is of two
types, emission and absorption
...
These radiations when analysed with the help of
spectroscope, spectral lines are obtained
...
Emission spectra is of two types,
(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 help of a spectroscope a series of thin bright lines of specific colours
are obtained
...
This type of
spectrum is called line spectrum or atomic spectrum
...
On analysing the transmitted light we obtain a spectrum in which dark lines of
specific wavelengths are observed
...
The
wavelength of the dark lines correspond to the wavelength of light absorbed
...

(2) When an electric discharge is passed through hydrogen gas at low pressure, a bluish
light is emitted
...

(4) All these lines of H-spectrum have Lyman, Balmer, Paschen, Barckett, Pfund and
Humphrey series
...

(5) To evaluate wavelength of various H-lines Ritz introduced the following expression,
1
c

R

1
n 12

1
2
n2

Where R is universal constant known as Rydberg‟s constant its value is 109, 678 cm

1


...

(1) Thomson regarded atom to be composed of positively charged protons and negatively
charged electrons
...

(2) He regarded the atom as a positively charged sphere in which
negative electrons are uniformly distributed like the seeds in a water
melon
...


–

–

+ – +
–

+
–

+

Electron

+

Positive charge spreaded throughout the sphere

Rutherford's nuclear model
...


(ii) Some of them were deflected away from their path
...

(iv) The scattering of

1

particles
sin


...
e
...

(ii) Extra nuclear part which contains electrons
...

Planetry electron
–
Nucleus

+
10–15 m

10

–10

m

Size of the nucleus = 1 Fermi = 10
Size of the atom 1 Å = 10

–10

–15

m

m

(3) Properties of the Nucleus
(i) Nucleus is a small, heavy, positively charged portion of the atom and located at the
centre of the atom
...
e
...


(iii) Nucleus contains neutrons and protons, and hence these particles collectively are also
referred to as nucleons
...

(v) The radius of nucleus is of the order of 1
...
to 6
...
i
...
1
...
5 Fermi
...


ro ( 1
...
e
...


(vii) The density of the nucleus is of the order of 10 15 g cm

3

or 10 8 tonnes cm

3

or

10 12 kg / cc
...
023 10 23
r
3

(4) Drawbacks of Rutherford's model
(i) It does not obey the Maxwell theory of electrodynamics, according to it “A
small charged particle moving around an oppositely charged centre continuously
loses its energy”
...

(ii) It could not explain the line spectra of H atom and discontinuous spectrum
nature
...


Planck's Quantum theory
(1) Max Planck (1900) to explain the phenomena of 'Black body radiation' and
'Photoelectric effect' gave quantum theory
...

(2) If the substance being heated is a black body (which is a perfect absorber and perfect
radiator of energy) the radiation emitted is called black body radiation
...
In case of light, the quantum of
energy is called a 'photon'
...
e
...
62×10–27 erg
...
or 6
...


(iii) The total amount of energy emitted or absorbed by a body will be some whole number
quanta
...

(iv) The greater the frequency (i
...
shorter the wavelength) the greater is the energy of
the radiation
...


Photoelectric effect
(1) When radiations with certain minimum frequency ( 0 ) strike the surface of a metal, the
electrons are ejected from the surface of the metal
...
The current constituted by
photoelectrons is known as photoelectric current
...
The minimum potential at which
the plate photoelectric current becomes zero is called stopping potential
...


(4) The number of photoelectrons ejected is proportional to the intensity of incident
radiation
...

U
...
light

Visible light

Note :  Nearly

Metal

Metal other
than alkali
metals

Photo electrons

No photo electrons

Alkali
allVisible light emit photoelectrons when exposed to u
...
light
...


and caesium

emit

 Cesium (Cs) with lowest ionization energy among alkali metals is used in photoelectric
cell
...

(1) This model was based on the quantum theory of radiation and the classical law of
physics
...

(2) Postulates of this theory are :
(i) The atom has a central massive core nucleus where all the protons and neutrons are
present
...

(ii) The electron in an atom revolve around the nucleus in certain discrete orbits
...

(iii) The force of attraction between the nucleus and the electron is equal to centrifugal
force of the moving electron
...
Thus, mvr n
2
Where, m = mass of the electron, r = radius of the electronic orbit, v = velocity of the
electron in its orbit
...


...
In the above equation „n‟ is any integer
which has been called as principal quantum number
...
Various energy levels are designed as K(n=1),
L(n=2), M(n=3) ------- etc
...


(v) The angular momentum can be

(vi) The emission or absorption of radiation by the atom takes place when an electron
jumps from one stationary orbit to another
...

Thus, h

E

Where „h‟ =Planck‟s constant,
frequency of the radiant energy
...

(viii) The lowest energy state (n=1) is called the ground state
...
It has to fall back to a
lower orbit with the release of energy
...

hydrogen atom, He , Li 2 etc
...

4 2 me 2 k Z

where, n =Orbit number, m =Mass number 9
...
6 10
9 10 9 Nm 2 c

19

31

kg , e =Charge on the

Z =Atomic number of element, k = Coulombic constant

2

After putting the values of m,e,k,h, we get
...
529 Å or rn

n2
Z

0
...
g
...
e
...
:: 1 : 4 : 9
...
188 10 8
cm
...
or V1 : V2 : V3
...

2 3

(b) For a particular orbit of different species

V Z [n =constant] Thus, we have H
(c) For H or He+ or Li+2, we have

V1 : V2

2 : 1 ; V1 : V3

3 : 1 ; V1 : V4

He

Li

2

4 :1

(iv) Calculation of energy of electron in Bohr’s orbit
Total energy of electron = K
...
+ P
...
of electron
Substituting of r, gives us E
Putting the value of m, e, k, h,
E

21
...
6

12

2

2

mZ 2 e 4 k 2

n2h2
we get

Z2
erg per atom
n2

Z2
eV per atom (1eV
n2

21
...
6 10 -19 J )

1312 2
Z kJmol 1
n2
(a) For a particular system[H, He+ or Li+2]
or

E
1
[Z =constant] Thus, we have 1
2
E2
n

kZe 2
r

kZe 2
2r

Where, n=1, 2, 3………
...
18J)

E

kZe 2
2r

2
n2
2
n1

The energy increase as the value of n increases
(b) For a particular orbit of different species

Z2
J per atom (1 J
n2
313
...
/ mole
n2

(1 cal =

E

For the system
H

He

Li

2

2
ZA

E
Z [n =constant] Thus, we have A
EB
2

2
ZB

H, He+ , Li+2, Be+3 (n-same) the energy order is

3

Be
The energy decreases as the value of atomic number Z increases
...
6 Z 2

E

1
n12

2

2

1
2
n2

1
eV / atom
2
n2

k 2 me 4 Z 2 1
2
ch 3
n1

where, R

2

2

1
2
n2

k 2 me
ch 3

4

R is

known as Rydberg constant
...

(4) Quantisation of energy of electron
(i) In ground state : No energy emission
...

E1
13
...

(ii)

In excited state : Energy levels greater than n1 are excited state
...
e
...
For H- atom first excitation state is

n2

(iii) Excitation potential : Energy required to excite electron from ground state to any
excited state
...
4 13
...
2 eV
...
5 13
...
1 eV
...

E ionisation

E

En

13
...

eV
...

S
...
= E

Eexcited
...
E
...
6

...
E
...
These spectral series were named by the name of scientist who
discovered them
...
581 cm

1

This remarkable agreement between the theoretical and experimental value was
great achievment of the Bohr model
...

n=8
n=7
n=6
n=5

Humphrey series
Pfund

Energy level

n=4

series
Brackett
series

n=3

Paschen

series

n=2
Balmer
series

n=1

Lyman
series

(iv) Comparative study of important spectral series of Hydrogen
S
...


Spectral
series

Lies
in
the region

Transitio
n

n2
(1)

(2)

(3)

Lymen
series

Balmer
series

Paschen
series

Ultraviolet
region

Visible
region

Infra
red
region

max

1

n2

2,3,4
...


max

n1
max

n1 = 3

n1

4,5,6
...


n1

5

n2

6,7,8
...


max

n1
max

n1
max

1 and n 2

2

4
3R

2 and n 2

min

3

36
5R

3 and n 2

4

5

36 49
13 R

n1
min

6

25 36
11 R

6 and n 2

n1
min

16 25
9R

5 and n 2

n1
min

144
7R

4 and n2

n1

n1
min

7

n1
min

1 and n 2
1
R

4
3

2 and n2
4
R

3 and n 2
9
R

9
5

16
7

4 and n 2
16
R

5 and n 2
25
R

6 and n 2
36
R

25
9

36
11

49
13

(v) If an electron from nth excited state comes to various energy states, the maximum
n(n 1)

...

spectral lines obtained will be =
2
6(6 1) 30
15
...

(6) Failure of Bohr Model
(i) Bohr theory was very successful in predicting and accounting the energies of line
spectra of hydrogen i
...
one electron system
...

(ii) This theory could not explain the presence of multiple spectral lines
...
The intensity of these spectral
lines was also not explained by the Bohr atomic model
...

(v) This theory could not explain uncertainty principle
...


Bohr – Sommerfeld’s model
...

(2) He gave concept that electron revolve round the nucleus in elliptical orbit
...

nh
(3) For circular orbit, the angular momentum =
where n= principal quantum number
2
only one component i
...
only angle changes
...
In elliptical orbit two
components are,
h
(i) Radial component (along the radius) = nr
Where, n r = radial quantum number
2
h
(ii) Azimuthal component = n
Where, n = azimuthal quantum number
2
h
h
n
So angular momentum of elliptical orbit = nr
2
2
Angular momentum = (n r

n )

h
2

r

2

2

rr

r

rr
1
1

1

(5) Shape of elliptical orbit depends on,

Length of major axis
Length of minor axis

n
n

nr

= change

= change

n

r = change

n

r = constant

(6) n can take all integral values from l to „n‟ values of n r depend on the value of n
...
e
...

Thus for n = 3, we have 3 paths
n

n

nr

Nature of path

3

1
2
3

3
1
0

elliptical
elliptical
circular

K= 3
K= 2
K= 1
Nuclear

The possible orbits for n = 3 are shown in figure
...
therefore it provides the basis for existance of subshells in Bohr's shells (orbits)
...

(ii) When electron jumps from inner orbit to outer orbit or vice –versa, then electron run
entire distance but absorption or emission of energy is discontinuous
...

2
This model could not explain Zeeman effect and Stark effect
...

(1) In 1924, the french physicist, Louis de Broglie suggested that if light has both particle
and wave like nature, the similar duality must be true for matter
...

(2) This presented a new wave mechanical theory of matter
...

(3) According to de-broglie, the wavelength associated with a particle of mass m, moving
with velocity v is given by the relation

h
, where h = Planck‟s constant
...
c

c

energy of photon (on the basis of Einstein‟s mass energy relationship), E
equating both

hc

mc 2 or

h
which is same as de-Broglie relation
...
Let the electron is accelerated with a potential of V than the
Kinetic energy is

1
mv 2
2

eV ; m 2 v 2

2eVm and mv

2 eVm

P;

h
2 eVm

(6) If Bohr‟s theory is associated with de-Broglie‟s equation then wave length of an
electron can be determined in bohr‟s orbit and relate it with circumference and multiply with a
2 r
2 r n or
whole number
n
From de-Broglie equation,

Note : 

h
h

...


(7) The de-Broglie equation is applicable to all material objects but it has significance only
in case of microscopic particles
...


Heisenberg’s uncertainty principle
...

(2) According to uncertainty principle “It is impossible to specify at any given moment both
the position and momentum (velocity) of an electron”
...
p

x
...
If x is made small, p increases and vice versa
...
t

h
4

: Heisenberg‟s uncertainty principle cannot we apply to a stationary electron

because its velocity is 0 and position can be measured accurately
...

(1) Schrodinger wave equation is given by Erwin Schrödinger in 1926 and based on
dual nature of electron
...

(3) The probability of finding an electron at any point around the nucleus can be
determined by the help of Schrodinger wave equation which is,
2

x

2
2

y

2
2

z

2

8
2

m

h

2

(E

V)

0

Where x, y and z are the 3 space co-ordinates, m = mass of electron, h = Planck‟s
constant,
E = Total energy, V = potential energy of electron,
as wave function
...


= amplitude of wave also called

(4) The Schrodinger wave equation can also be written as :
2

8

2

m

h

Where

2

(E

V)

0

= laplacian operator
...
The
amplitude
is thus a function of space
co-ordinates and time i
...

(x, y, z
...

(iii) If

2

is maximum than probability of finding e is maximum around nucleus
...
It is different from the Bohr‟s orbit
...


Quantum numbers and Shapes of orbital’s
...

(2) Principle quantum number (n)
(i) It was proposed by Bohr’s and denoted by „n‟
...

r

n2
Z

0
...

E

Z2
n2

313
...
No energy shell in atoms of known elements possess more than 32
electrons
...

(vi) The value of energy increases with the increasing value of n
...

(viii) Angular momentum can also be calculated using principle quantum number

nh
2

mvr
(3) Azimuthal quantum number (l)

(i) Azimuthal quantum number is also known as angular quantum number
...

(ii) It determines the number of sub shells or sublevels to which the electron belongs
...

(iv) It also expresses the energies of subshells s
(v) The value of l

p

f (increasing energy)
...


(vi) Value of l

=

0

1

2

3………
...
Which is equal to
(viii) The maximum number of electrons in subshell

s

h
2

l(l 1)

2(2l 1)

subshell

2 electrons d

subshell

10 electrons

p subshell

6 electrons f

subshell

14 electrons
...

(x) The energy of any electron is depend on the value of n & l because total energy = (n
+ l)
...

(4) Magnetic quantum number (m)
(i) It was proposed by Zeeman and denoted by „m‟
...

(iii) The value of m varies from –l to +l through zero
...
e
...

(v) For a given value of „n‟ the total value of ‟m‟ is equal to n 2
...

(vii) Degenerate orbitals : Orbitals having the same energy are known as degenerate
orbitals
...
g
...

(5) Spin quantum numbers (s)
(i) It was proposed by Goldshmidt & Ulen Back and denoted by the symbol of „s‟
...


(iii) The spin may be clockwise or anticlockwise
...


(iv) It represents the value of spin angular momentum is equal to
(v) Maximum spin of an atom

s(s

1)
...


Shape of orbitals
(1) Shape of ‘s’ orbital
(i) For „s‟ orbital l=0 & m=0 so „s‟ orbital have only one
unidirectional orientation i
...
the probability of finding
the electrons is same in all directions
...


X

(iii)It does not possess any directional property
...


Nucleus

(2) Shape of ‘p’ orbitals
(i) For „p‟ orbital l=1, & m=+1,0,–1 means there are three „p‟ orbitals, which is
symbolised as p x , p y , p z
...


(iii) p-orbital has directional properties
...
It shows that the „d‟
orbitals has five
orbitals as d xy , d yz , d zx , d x 2 y 2 , d z 2
...

(iii) The shape of d orbital is double dumb bell
...

Z

Z

Y

Z

Y

Y

Y

Y
Z

X

X
dXY

dZX

X

X
dX2–Y2

dYZ

X
dZ2

(4) Shape of ‘f’ orbital
(i) For the „f‟ orbital l=3 then the values of „m‟ are –3, –2, –1,0,+1,+2,+3
...

(ii) The „f‟ orbital is complicated in shape
...

The distribution of electrons in different orbitals of atom is known as electronic
configuration of the atoms
...

(ii) According to this principle, “In the ground state, the atomic orbitals are filled in
order of increasing energies i
...
in the ground state the electrons first occupy the
lowest energy orbitals available”
...

(iv) According to this rule
(a) Lower the value of n + l, lower is the energy of the orbital and such an orbital
will be filled up first
...

Thus, order of filling up of orbitals is as follows:

1s

2s

2p

3s

3p

4s

4p

5s

4d

5p

6s

6f

5d

(2) Pauli’s exclusion principle
(i) According to this principle, “No two electrons in an atom can have same set of all
the four quantum numbers n, l, m and s
...

(iii)Since this principle excludes certain possible combinations of quantum numbers
for any two electrons in an atom, it was given the name exclusion principle
...

(b) The maximum capacity of a subshell is equal to 2(2l+1) electron
...

(d) Number of orbitals in a main energy shell is equal to n 2
...

(iv) According to this principle an orbital can accomodate at the most two electrons
with spins opposite to each other
...

(v) If an orbital has two electrons they must be of opposite spin
...

(ii) According to this rule “Electron filling will not take place in orbitals of same energy
until all the available orbitals of a given subshell contain one electron each with
parallel spin”
...

(iv) The reason behind this rule is related to repulsion between identical charged
electron present in the same orbital
...

(vi) Moreover, according to this principle, the electron entering the different orbitals of
subshell have parallel spins
...

(vii) The term maximum multiplicity means that the total spin of unpaired e is
maximum in case of correct filling of orbitals as per this rule
...


One electron system : In this system 1s 2 level and all orbital of same principal quantum
number have same energy, which is independent of (l)
...

Multiple electron system : The energy levels of such system not only depend upon the
nuclear charge but also upon the another electron present in them
...
For example
energy level of 2 > 1
...
For a
definite shell, the subshell having higher value of l possesses higher energy level
...

Energy level order

4f
l= 3

>

4d

>

4p

>

l=2

4s

l=1

l= 0

(iii) The relative energy of sub shells of different energy shell can be explained in the terms of the
(n+l) rule
...

For

3d

n=3

l= 2

n+l=5

For

4s

n=4

l=0

n+l=4

(b) If the value of (n + l) for two orbitals is same, one with lower values of ‘n’ possess lower
energy level
...

(b) Thus, if the shift of an electron from one orbital to another orbital differing slightly
in energy results in the symmetrical electronic configuration
...


(c) For example p 3 , d 5 , f 7 configurations are more stable than their near ones
...

(b) The energy is released during the exchange process with in the same subshell
...
g
...

(d) The greater the number of possible exchanges between the electrons of parallel
spins present in the degenerate orbitals, the higher would be the amount of energy
released and more will be the stability
...

d4
(2)
(1)
3 exchanges by 1st e–
2 exchanges by 2nd e–
To number of possible exchanges = 3 + 2 + 1 =6
d5
(1)

(2)
4 exchanges by 1st

(3)
Only 1 exchange by 3rd e–

(3)
3 exchanges by 2nd e–

e–

2 exchange by 3rd e–

(4)
1 exchange by 4th e–

To number of possible exchanges = 4 + 3 + 2 +1 = 10

Electronic configurations of Elements
...
C
...
Hence,
usually the electronic configuration of the atom of any element is simply represented by the notation
...
g
...


(3) (i) Elements with atomic number 24(Cr), 42(Mo) and 74(W) have ns 1 (n 1) d 5
configuration and not ns 2 (n 1) d 4 due to extra stability of these atoms
...

Cr (24)

3d5

[Ar]

Cu (29)

3d10

[Ar]
4s1

4s1

(4) In the formation of ion, electrons of the outer most orbit are lost
...
If we write electronic configuration of Fe
2

26 , 24 e ), it will not be similar to Cr (with 24 e ) but quite different
...

(5) Ion/atom will be paramagnetic if there are unpaired electrons
...
(1BM
only) is
unpaired electrons
...
27 10

24

J / T ) where n is the number of

(6) Ion with unpaired electron in d or f orbital will be coloured
...

10

2

with electronic

with electronic configuration Ar 3d 9 (one

(7) Position of the element in periodic table on the basis of electronic configuration can be
determined as,
(i) If last electron enters into s-subshell, p-subshell, penultimate d-subshell and anti
penultimate f-subshell then the element belongs to s, p, d and f – block respectively
...

(iii)If the last shell contains 1 or 2 electrons (i
...
for s-block elements having the
configuration ns 1 2 ), the group number is 1 in the first case and 2 in the second
case
...
e
...

(v) If the electrons are present in the (n –1)d orbital in addition to those in the ns
orbital (i
...
for d-block elements having the configuration (n –1) d 1 10 ns 1 2 ), the
group number is equal to the total number of electrons present in the (n –1)d
orbital and ns orbital

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