Structure of Atom

 According to Dalton’s atomic Theory, atom is

indivisible.
(Atom term is derived from Greek word a-tomio = uncut
able)
 Later scientists like J.J.Thomson, Goldstein,
Rutherford, Chadwick, Bohr and others established that
atom is divisible & made up of still smaller subatomic
particles like electron, proton, neutrons etc.
 At present 35 subatomic particles are known. e.g.
Higgs Boson, Quark, muon, graviton, neutrino, etc.
 But three particles electron, proton and neutron are
regarded as fundamental particles.
Discovery of Fundamental Particles:
 Discovery of Electrons:
(Discharge Tube Experiment)
Cathode rays were analyzed & it was found that cathode
rays are made up of small-small negatively charged
particles, called Electrons.
 𝒆
 J.J.Thomson (1897) determined ratio for electrons
𝒎
by his experiment.
𝒆 8
=1.76 * 10 coulomb/g
𝒎
𝒆
& it was found that ratio was same for all gases taken
𝒎
in discharge tube.

𝑻𝒉𝒊𝒔 𝒑𝒓𝒐𝒗𝒆𝒔 𝒕𝒉𝒂𝒕 𝒆𝒍𝒆𝒄𝒕𝒓𝒐𝒏 𝒊𝒔 𝒇𝒖𝒏𝒅𝒂𝒎𝒆𝒏𝒕𝒂𝒍 𝒑𝒂𝒓𝒕𝒊𝒄𝒍𝒆

  R.A.Millikan (1917) determined charge on electron
with the help of Oil Drop Experiment.
Charge of electron= 1.60 * 10-19 coulomb / 4.8 * 10-10 esu
 Mass of electron can be calculated as:
𝒆 𝟏.𝟔𝟎∗ 𝟏𝟎−𝟏𝟗
Mass = 𝒆 = 9.11 * 10-28 g.
𝟏.𝟕𝟔∗ 𝟏𝟎𝟖
𝒎
𝟏
Which is of mass of hydrogen atom.
𝟏𝟖𝟑𝟕
Thus electron is fundamental particle which carries one
𝟏
unit negative charge & has a mass equal to of that
𝟏𝟖𝟑𝟕
hydrogen atom.
Origin of Cathode rays:
Initially from the material of cathode & then from the
ionization of gas taken in discharge tube.
Discovery of Proton:
Goldstein (1886) pointed out that like cathode rays, anode
rays or canal rays also comes from anode side. (They
passed through canals i.e. holes in cathode & produces
fluorescence on back side of cathode)

Origin of Anode Rays:

Anode rays are produced in space between anode &
cathode by ionization of gas taken.
(Caused by bombardment of cathode rays over gaseous atom taken in discharge tube.)
 On analyzing anode rays it was found that they
remains consists of positively charged particles.
𝒆
 The value for these particles was found different
𝒎
for different gases.
 When Hydrogen gas is taken, lightest positively
charged particle is produced, called Proton.
 Charge of proton = 1.6 * 10-19 coulomb.
𝒆 4
 for proton = 9.58 * 10 coulomb/g
𝒎
𝒆 𝟏.𝟔∗ 𝟏𝟎−𝟏𝟗
 Mass of proton = 𝒆 = = 1.6726 * 10-24 g
𝟗.𝟓𝟖∗ 𝟏𝟎𝟒
𝒎
Thus, Proton is fundamental particle which carries one
unit positive charge & has mass equal to that of hydrogen
atom.
Discovery of Neutron:
Since mass of electron is negligible,
Therefore, it is expected:
mass of an atom = total mass of protons atom have,
Or Total number of protons atom have.
(as in atomic mass scale, mass of proton is taken unity.)
But, mass of an atom >> Total number of protons/total mass of protons.
Since overall atom is neutral, thus it was predicted the presence of some
other neutral particle in an atom.
Chadwick (1932) obtained neutral particle by his scattering
experiment, called Neutron.
𝟗 𝟒 𝟏𝟐 𝟏
𝟒 𝑩𝒆 + 𝟐 𝑯𝒆 → 𝟔 𝑪 + 𝟎𝒏
𝟏𝟏 𝟒 𝟏𝟒 𝟏
𝟓 𝑩 + 𝟐 𝑯𝒆 → 𝟕 𝑵 + 𝟎𝒏
-24
Mass of neutron = 1.6749 * 10 g
Neutron is fundamental particle present in an atom which carries a mass
1.675 * 10-24 g and no charge.
Discovery of Nucleus:
(Rutherford Scattering Experiment)

Rutherford bombarded α-particles (i.e. nucleus of Helium)

over thin foil (approx. 100 nm thick) of metals like gold,
silver, Pb or Cu.
Source of α-particles was “Ra” placed in block of lead.
Observations:
99.9% α-particles passed through foil without undergoing deflection.
Few particles show deflection through small angle.
Very few (one out of 20000) rebounds.

Conclusions:
 Maximum part of an atom is empty.
 Some positive charge is present within the atom.
 Entire mass of an atom is concentrated in a very
small region at the center of atom, called Nucleus.

From X-ray experiments, it is found:

-10
Radius of an atom = 10 m
-15
Radius of Nucleus = 10 m
Atomic Number: (i.e. Z)
Mosely (1913) analyzed the X-rays obtained by
bombardment of cathode rays over different elements &
found that frequency of emitted X-rays depends over
magnitude of positive charge of nucleus, as:
𝝂 = a(Z-b)
Here “a” & “b” are constants.
Since positive charge of nucleus is due to protons, thus:
Mosely states that number of positive charge present in
nucleus of an atom is known as atomic number of element.
Since, proton contains unit positive charge, Therefore:
Atomic number = Number of protons = Number of electrons
(As overall atom is neutral)
Mass Number: (i.e. A)
As mass of electron is negligible, Therefore entire mass of
an atom is due to mass of protons & neutrons.
(since proton & neutron has unit mass)
Therefore,
Mass Number = Number of protons + Number of Neutrons
 Mass number of an element remains nearly equal to
atomic mass of the element.
Mass Number Atomic Mass
Is always whole number. Usually not whole number.
 Pattern of Representation:
𝑨
𝒁 𝑿 Symbol
Isotopes:
Atoms of same element having same atomic number but different mass
numbers are called isotopes.
e.g. 𝟏 𝟐
𝟏𝑯 𝟏𝑯 & 𝟑𝟏𝑯, 𝟏𝟔 𝟏𝟔
𝟖𝑶 𝟖𝑶 & 𝟏𝟖
𝟏𝑶 , 𝟑𝟓
𝟏𝟕𝑪𝒍 & 𝟑𝟕
𝟏𝟕𝑪𝒍 , 𝟏𝟐 𝟏𝟐
𝟔𝑪 𝟔 𝑪 & 𝟏𝟐
𝟔𝑪

Isobars:
Atoms of different elements having same mass numbers but different
atomic numbers are called isobars.
𝟒𝟎 𝟒𝟎 𝟒𝟎
e.g. 𝟏𝟖𝑨𝒓 𝟏𝟗𝑲 & 𝟐𝟎𝑪𝒂

Isotones:
Atoms of different elements having same number of neutrons are called
isotones.
e.g. 𝟏𝟒𝟔𝑪 𝟏𝟓𝟕𝑵 & 𝟏𝟔𝟖𝑶, These all have 8 neutrons in their nucleus.
Isoelectronics:
Species (atoms or ions) containing same number of electrons are called
isoelectronics.
e.g. 𝑵𝟑− , 𝑶𝟐− , 𝑭− , 𝑵𝒂+ , 𝑴𝒈𝟐+ , 𝑨𝒍𝟑+ , & Ne.
Each of them contains 10 electrons.
Electromagnetic Wave Theory:
(James Clark Maxwell 1864)
The energy is emitted from source continuously in the form of
radiation.
i.e. radiations of only one wave length given out, irrespective of its
temperature.
Therefore, Energy of radiations α intensity of radiations
Radiation consists of electric field & magnetic field oscillating
perpendicular to each other as well as to the direction of propagation .
Thus called Electromagnetic radiation/wave.
All electromagnetic radiations travels with velocity of light i.e. 3 * 108
m/s
Electromagnetic radiations do not require any medium for propagation.
Characteristics of Wave:
 Wavelength(i.e. 𝝀):distance between to consecutive crests or trough.
 Frequency(i.e.𝝂):number of waves passing through a point in one second.
 Amplitude (i.e. a):maximum height of crest or depth of the trough.
 Wave Number (i.e. 𝝂):number of waves present in 1 cm length.
𝟏
𝝂=
𝝀
Relation between Velocity, 𝝂 & 𝝀 :
C=𝝂*𝝀
Electromagnetic Spectrum:
i.e. Arrangement of different electromagnetic radiations in
order of increasing wavelength or decreasing frequencies.

VIB G Y O R
Limitations of Electromagnetic Wave Theory:

Theory has successfully explained some

properties of light such as interference,
diffraction, etc. but could not explain:

 Phenomenon of black body radiation.

 Photoelectric effect.
 Line spectra of atoms. (i.e. of Hydrogen)
Planck’s Quantum Theory of Radiations:
(Max Planck 1900)
 Radiant energy is emitted discontinuously in the
form of small energy packets, called quanta. (In case of
light these are called photon.)
 The energy of each quanta is directly proportional to
the frequency of associated radiation.
𝒄
E α 𝝂, E = h 𝝂, E = h
𝝀
h = Planck’s Constant.(6.626*10-34 js/3.99*10-13kj sec mol-1)
 Total energy absorbed or emitted is integral multiple
of quantum.
𝒄
E = nh 𝝂 = nh (Here n= 1,2,3,----etc.)
𝝀
Energy of 1 mole of quanta is called Einstein
of energy.
Black body radiation:
An ideal body which emits and absorbs radiations of all
wavelengths or frequencies is called black body & the
radiations emitted by black body is called black body
radiation.
When a black body, such as iron rod is heated in furnace,it
undergoes series of changes in colour.
Dull red → more red → yellow → white → blue

According to wave theory, black body is not expected to

show change in colour.
Explanation by Planck’s quantum theory:
According to this theory energy emitted from source
discontinuously, in the form of radiation consisting of small
energy packets called Quanta.
The energy of each quantum remains proportional to
frequency of radiation associated.
Eα𝝂
Thus when a black body is heated, with increase in
temperature, the energy & therefore frequency of emitted
radiation also changes.
Since there is co-relation between frequency & colour,
therefore change in colour is justified .
Photoelectric Effect:
The phenomenon of ejection of electron from the surface
of a metal, when light of suitable frequency strikes on it is
called photoelectric effect & emitted electrons are called
Photoelectrons.

 One photon ejects one electron. Therefore number

of electrons ejected per second depends upon intensity
of incident radiation & does not depend over its
frequency.
Explanation:
 Light consists of stream of Photons.
 Striking photons must have energy higher than
threshold energy (i.e. h𝝂𝟎 ).
 Threshold energy is also called work function.
 The extra energy of striking photon turns to kinetic
energy of ejected electron.
Energy of = threshold + kinetic energy of
Incident photon energy ejected electron
𝟎 𝟏
(i.e. h𝝂) (i.e. h𝝂 ) ( 𝒎𝒗𝟐 )
𝟐
𝟏
𝒎𝒗𝟐 = h𝝂 - h𝝂𝟎
𝟐
𝟏 𝟐 𝟎
𝒎𝒗 = h (𝝂 - 𝝂 )
𝟐
 Spectrum
Discontinuous Spectrum
Continuous Spectrum
Or
 Here there is Line Spectrum
continuity of colours  Consists of lines of
i.e. one colour definite frequencies.
merges into other  e.g. atomic spectrum
without any gap.
 e.g. rainbow
Atomic Spectrum:
Consists of sharp well defined lines or bands
corresponding to definite frequencies. These are of two
types.
(1)Emission spectra (2) Absorption spectra.
Emission Spectra:
Obtained when radiation emitted from source are directly
analyzed with the help of spectroscope.

Source Spectroscope
Radiations can be obtained ;
From sun or glowing electric bulb
By passing electric discharge through a gas at low pressure.
By heating substance at high temperature.
Absorption Spectra:
Obtained when radiations from source is first allowed to
pass through sample & then the transmitted light is
analyzed by the spectroscope.

Source Sample Spectroscope

Absorption spectrum is always discontinuous, i.e. has few
dark lines due to absorption of light by sample.
 It is noted that for an element , dark lines obtained exactly at the
same place, where the lines are obtained for the element in its
emission spectra.
e.g. emission spectra of sodium consists of two yellow lines at 589 nm &
589.6 nm. When light is passed through sodium vapour, we obtains two
dark lines at 589 & 589.6 nm.
Rutherford Nuclear Model of Atom:
 In atom, entire mass & positive charge is
concentrated in very small region at centre called
nucleus.
 Positive charge of nucleus is due to protons.
 Nucleus remains surrounded by negatively charged
electrons, & number of protons = number of electrons.
 Electrons revolves round the nucleus in circular path
with very high speed.
 The mass of nucleus is due to protons & some other
neutral particle having mass almost equal to that of
proton. (later Chadwick discovered it & named
Neutron)
 Most of the space of an atom between nucleus &
revolving electrons is empty.
Drawbacks of Rutherford model:
 Fails to explain the stability of an atom.
(acc. To Clark Maxwell, a charged body moving under the influence of
attractive forceloses energy in the form of electromagnetic radiations.
Consquently electron will come closer & closer to nucleus & finally will
fell in nucleus, i.e. atom will collapse. But it never happens.)

Fails to explain the line spectrum of Hydrogen.

(Acc. To Rutherford model, Hydrogen spectrum must be continuous.)
Bohr’s Atomic Model: (Neils Bohr : 1913)
Atom consists of small heavy positively charged nucleus at centre
surrounded by electrons.
Electrons revolves round the nucleus in fixed energy circular path called
orbits.
Each orbit has fixed radius & energy.
For Hydrogen atom,
𝟐𝝅𝟐 𝒎𝒆𝟒 𝟐𝟏.𝟖∗ 𝟏𝟎−𝟏𝟗 𝟏𝟑.𝟔 𝟏𝟑𝟏𝟐
𝑬𝒏 = - =- J/atom = - eV/atom = - kj/mol
𝒏𝟐 𝒉𝟐 𝒏𝟐 𝒏𝟐 𝒏𝟐
For Hydrogen like ions i.e. 𝑯𝒆+ , 𝑳𝒊 𝟐+
, etc
𝟐𝝅𝟐 𝒎𝒆𝟒 ∗𝒁𝟐 𝟏𝟑𝟏𝟐∗𝒁𝟐
𝑬𝒏 = - =- kj/mol
𝒏𝟐 𝒉𝟐 𝒏𝟐
𝟐
𝒓𝒏 = 𝒂𝟎 𝒏 (𝒂𝟎 =52.9 pm) For Hydrogen
For Hydrogen like ions i.e. 𝑯𝒆+ , 𝑳𝒊𝟐+, etc
𝒂𝟎 𝒏𝟐
𝒓𝒏 =
𝒁
𝒁
Velocity of electron, 𝒗𝒏 = 𝒗𝟎
𝒏
Here 𝒗𝟎 = 2.188 * 108 cms (velocity of electron in 1st orbit of H atom)
-1

(energy of orbit increases with increase in distance from nucleus.)

  Infinite number of orbits are possible.
 Electron revolves in only those orbits for which
angular momentum of electron is whole number
𝒉 𝒉
multiple to . i.e. mvr = n *
𝟐𝝅 𝟐𝝅
 Electron does not lose or gain energy, till they
revolves in its orbit.
(it loses or gains energy, only when they changes their orbit.)
Achievements of Bohr’s Model:
(1) Successfully explained the stability of atom.
(Question of losing energy continuously does not
arise, as electron does not lose energy till it revolves
in its own orbit.)
(2) Successfully explained the line spectrum of Hydrogen

Emission Spectrum of Hydrogen Atom:

When electric discharge is passed to sample of hydrogen molecules,
these splits into atoms. Now different hydrogen atom absorbs different
amount of energy, & therefore electron excites to different different
energy levels. As excited state is unstable, thus unstable electron come
back by losing energy. Energy is lost in the form of electromagnetic
radiation. These radiations when falls over photographic film, leaves its
impression as line.
  Simultaneous appearance of large number of lines in
hydrogen spectrum, although it has only one electron,
is because we are taking sample of H2, which contains
large number of molecules & therefore large number of
atoms.
 Number of emission lines obtained, when electron
jumps from 𝒏𝟐 to 𝒏𝟏 :
𝒏𝟐 +𝒏 𝟏 (𝒏𝟐 −𝒏𝟏 + 𝟏)
=
𝟐
Rydberg Equation:
𝟏 𝟏 𝟏
= 𝝂 = 𝑹𝑯 ( - )
𝝀 𝒏𝟐𝟏 𝒏𝟐𝟐
+ 2+
For Hydrogen like ions i.e. He , Li etc.
𝟏 𝟐 𝟏 𝟏 -1
= 𝝂 = 𝑹𝑯 * 𝒁 ( 𝟐 - 𝟐) Here 𝑹𝑯 = 109678 cm
𝝀 𝒏𝟏 𝒏𝟐
Limitations of Bohr’s Model:

 Inability to explain line spectrum of multi-electron

atom.
 Inability to explain splitting of spectral lines in
magnetic field (i.e. Zeeman Effect) & in electric field (i.e.
Stark Effect).
 Inability to explain the shape of molecules.
 Inability to explain dual nature of matter &
Heisenberg uncertainty principle.
 Inability to explain three dimensional model of
atom.
(Bohr gave flat model for atom.)
Dual Behaviour of Matter:
De-Broglie Equation:
All matter particles in motion possess wave
characteristics.
According to De-Broglie, the wave length
associated with particle of mass m moving
with velocity v:
𝒉
𝝀=
𝒎𝒗
The wave associated with matter particle is
called matter wave or De-Broglie wave.
Derivation:
According to Planck’s quantum theory, energy of each
energy packet (i.e. quantum), E = h𝝂 ---------------(1)
According to Einstein, energy of each energy packet (i.e.
photon), E = m𝒄𝟐 -----------------------(2)
From equation (1) & (2)
𝟐 𝒄
h𝝂 = m𝒄 , as 𝝂 =
𝝀
𝒄
therefore, h = m𝒄𝟐
𝝀
𝒉𝒄 𝒉
𝝀= 𝟐=
𝒎𝒄 𝒎𝒄
Here c = velocity of light, if particle is moving with velocity
v then :
𝒉
𝝀=
𝒎𝒗
Significance of De-Broglie Equation:

Although De-Broglie equation is applicable to

all material objects but it has significance
only in case of microscopic particles.

For macroscopic particles 𝝀 is too small, that

cannot be measured by any available
method.
𝟏 𝟏
As 𝝀 α or 𝝀 α
𝒎𝒗 𝒑
Momentum of moving particle.
Heisenberg Uncertainty Principle:
( Werner Heisenberg, 1927)
It is not possible to measure simultaneously
the position & the momentum (or velocity) of
a microscopic particle with absolute accuracy.
Methametical Expression:
𝒉 𝒉
𝚫𝒙 * 𝚫𝒑 ≥ OR 𝚫𝒙 * 𝒎𝚫𝒗 =
𝟒𝝅 𝟒𝝅
Equality sign denotes minimum uncertainty.
𝚫𝒙 = uncertainty in measuring position.
𝒎𝚫𝒗 = uncertainty in measuring momentum
(or velocity)
Reason:
We cannot see any microscopic particle without
disturbing it.

When a beam of light reflected from surface of object reaches to our eye, then objects
becomes visible. In case of microscopic particle, due to collision of incident photon
microscopic particle gets deflected from their normal path & both direction as well as
velocity gets changed.
Heisenberg Uncertainty Principle is not
applicable over macroscopic particles, as
photons are very small & therefore will not
deflect macroscopic particle from their
normal path.

Thus position & velocity of moving

macroscopic particle can be measured with
100% accuracy.
Shape of Atomic Orbitals:
Orbital is region of space around the nucleus in which the
probability of finding electron is maximum (about
90%).Probability at any point around the nucleus is
calculated using Schrodinger wave equation & is
represented by density of points. The shape of electron
cloud thus obtained gives the shape of the orbital.
Shape of s-orbital:
s-orbital is spherical in shape.
Shape of p-orbital:
p-orbital is Dumb-bell shaped.
Shape of d-orbital:
d-orbital has double dumb-bell
𝒅𝟐𝒛 has a doughnut shaped electron cloud, while
others have clover leaf shape.

𝒅𝟐𝒛 has no nodal plane.

Shape of f-orbital:
f-orbitals has diffused shape & has 6 lobes.
Quantum Numbers:
Defined as a set of four numbers with the help of which we
get complete information about all electrons in an atom,
i.e. energy, orbital they belongs & also their spin.
(These numbers are like postal address of electrons.)
(1) Principal Quantum Number: Denoted by “n” & gives
information regarding the orbit/shell to which
electron belongs & also about energy associated to it.
This also gives information about average distance of
electron from the nucleus.
N = 1,2,3,4, -------etc.
For hydrogen & hydrogen like ions i.e. He+, Li2+ etc.
𝟐𝝅𝟐 𝒎𝒆𝟒 𝒛𝟐
𝑬𝒏 = -
𝒏𝟐 𝒉𝟐
(2) Azimuthal/Subsidiary/Secondary/Angular
Quantum Number: Denoted by “l” and gives
information regarding sub-orbit to which electron belongs
and also gives information about the angular momentum
of electron. This also gives information regarding shape of
sub-orbits & orbitals.
𝒉
Angular momentum of electron = 𝒍(𝒍 + 𝟏)
𝟐𝝅
Value of “l” depends over value of “n”
L = 0 to (n-1)
Let n = 4
l= 0 1 2 3
Designation s p d f
of sub-orbit
(3) Magnetic Quantum Number: Denoted by “m”
/ “𝒎𝒍 ” & gives information about the orbital to which
electron belongs.
(An electron because of its angular motion around the nucleus produces an electric field,
which in turn produces magnetic field. Now magnetic field of electron interacts with
magnetic field of earth. As a result electron present in particular sub-shell acquires certain
specific orientation in space around nucleus called orbitals. Each orientation corresponds to
an orbital. Therefore “m” determines the number of orbital in any sub-orbit.)
 Value of “m” depends over “l”
For any value of “l”, 𝒎𝒍 = -l to +l including 0.
Value of “l” Name of Value of 𝒎𝒍 Name
sub-shell of
orbital
l =0 s-suborbit 0 s-orbital
l=1 p- suborbit -1, 0, +1 p-
l=2 d- suborbit -2,-1,0,+1,+2 d-
l=3 f- suborbit -3,-2,-1,0,+1,+2,+3, f-
Maximum number of orbital in any sub-orbit = (2l+1)
(4) Spin Quantum Number:
Denoted by “s” / “𝒎𝒔 ” .

This Quantum number gives information regarding the spin

of electron around its axis.

As electron can spin either clockwise or anti-clockwise,

therefore spin quantum number can have maximum two
𝟏 𝟏
values i.e. + & - .
𝟐 𝟐

Spin quantum number leads spin angular momentum i.e.

𝒉
𝒎𝒔 = 𝒔(𝒔 + 𝟏)
𝟐𝝅
Pauli’s Exclusion Principle:
An orbital can have maximum two electrons & they must
have opposite spin.
OR
For the two electrons of same orbital, the value of all the
four quantum numbers cannot be same.

e.g. Let configuration is 2s2 n l 𝒎𝒍 𝒎𝒔

For 1st electron 2 0 0 +
𝟏
𝟐
For 2nd electron 2 0 0 -
𝟏
𝟐
Aufbau Principle:
In the ground state of the atoms, the orbitals are filled in
order of their increasing energies i.e. electrons first occupy
the lowest energy orbital available & inter into higher
energy orbitals only when the lower energy orbitals are
filled.
Bohr-Bury’s Rule/(n+l) Rule:
Energy of orbitals is determined by the Quantum Numbers
“n” & “l”.
 Orbitals fills in the order of increasing value of (n+l).
e.g.
Orbital 3d 4s
(n+l) value 3+2=5 4+0=4
4s orbital fills first, as it has lower value of (n+l).
 If the two orbitals have same value of (n+l), then the
orbital with lower value of “n” will be filled first.
Orbitals (n+l) value
4p 4+1=5
3d 3+2=5
Here 3d orbital fill first, as it has lower value of “n”.
Hund’s Rule of Maximum Multiplicity:
Pairing of electrons in the orbitals of same suborbit takes
place only when all the orbitals of that particular suborbit
gets singly filled with parallel spin.
e.g.
1
3p ↑
2
3p ↑ ↑
3
3p ↑ ↑ ↑
4
3p ↑↓ ↑ ↑
Here
th
4 electron will cause pairing.
Electronic Configuration:
Distribution of electrons into different shells, subshells &
orbitals of an atom is called its electronic configuration.
The electronic configuration of any orbital can be simply
represented as: Number of electrons

𝒏𝒍𝒙 symbol of sub-shell or orbital

Number of Shell

Stability of Fully Filled & Half Filled Orbitals:

Half filled & fully filled configurations acquires extra
stability, due to following reasons:
(a) Symmetrical distribution of electrons.
(symmetry leads stability)
(b) High exchange energy of stabilization.