1. The document discusses principles of quantum chemistry including classical mechanics and its inadequacies in explaining phenomena at the atomic level, Planck's quantum theory, and properties of electromagnetic radiation. 2. Key concepts covered include de Broglie's equation describing the wave-like nature of matter, Heisenberg's uncertainty principle, explanations of photoelectric effect and blackbody radiation. 3. The document also introduces quantum numbers, Hund's rule, Pauli's exclusion principle, and Aufbau's principle, which describe allowable electron configurations in atoms and molecules.

1. QUANTUM CHEMISTRY - 01
Principles of Quantum Chemistry
&
Concept of Quantum Numbers
&
Hund’s, Pauli’s & Auf-Bau’s Principle
Prepared by:
Prof Sarala Prasanna Pattanaik
Assistant Professor
Department of Chemistry

2. Magic of Quantum Chemistry

3. Magic of Quantum Chemistry

4. Current Slide Contains the topic on:
 Classical Mechanics and it’s inadequacies
 Concept of Planck’s Quantum theory
 Properties of Electromagnetic radiation
 Dual nature of matter (De-Broglie’s equation)
 Heisenberg’s uncertainty principle
 Photoelectric effect
 Blackbody radiation and related laws
 Concept of Quantum Numbers
 Hund’s Rule & Pauli’s Exclusion Principle
 Auf-Bau’s Principle or Building up Principle

5. Classical Mechanics
• The development of Classical mechanics is based on Sir Issac
Newton’s three laws of motion such as Law of Inertia, Law of
Force or momentum, Law of action and reaction.
• These laws are very useful to derive the relationship between
velocity, acceleration, momentum, force, work and energy etc.
• Classical mechanics explain correctly the motion of macroscopic
bodies such as planets, stars, pendulum, projectiles etc under the
influence of forces or with the equilibrium of bodies when all
forces are balanced. The subject may be thought of as the
elaboration and application of basic postulates first enunciated by
Isaac Newton in his Laws of motion.
• But fails to explain when applied to microscopic particles such as
atoms, molecules, nucleus, electrons etc.

6. Newton’s 1st Law of Motion
The first law of motion states that an object either continues to
remains at rest or continues its state of motion at a constant
velocity in a straight line, unless it is acted upon by an external
unbalanced force. This law is also known as Law of Inertia.

7. Newton’s 2nd Law of Motion
The 2nd law of motion states that the rate of change of momentum of
an object is directly proportional to the force applied or for an object
with constant mass, that the net force on an object is equal to the
mass of that object multiplied by the acceleration. This law is also
known as Law of force or momentum or acceleration.
For a body/particle of mass m, force can be written in the form
F = ma = mv/t = p/t, where F is force, a is acceleration, v is velocity
and p = mv is the momentum of the particle.

8. Newton’s 3rd Law of Motion
The 3rd law of motion states that when one object exerts a force on a
second object, that second object also exerts a force that is equal in
magnitude and opposite in direction on the first object i. e for every
action there is an equal and opposite reaction. This law is also known
as Law of action and reaction or the Law of opposing forces.

9. Inadequacies of Classical Mechanics
The inadequacies of Classical Mechanics are as given below.
• It does not hold good in the region of atomic dimensions.
• It fails to explain the dual nature of electromagnetic radiations.
• It could not explain the observed spectra of black body radiation.
• The variation of the specific heat of solids with temperature is not
explained.
• It could not explain the origin of the discrete spectra of atoms like
hydrogen. This is so because according to classical mechanics the
energy changes are always continuous.
• It fails to explain the photoelectric effect.
Failure of classical mechanics leads to the development of Quantum
mechanics which successfully explains the dual nature, blackbody
radiation, photoelectric effect, concept of uncertainty in determining
the position and momentum of microscopic particles simultaneously,
concept of orbit and orbital with the help of Schrodinger wave
equation in terms of wave function (ψ).

10. Planck’s Quantum Theory
• Max Planck in 1900 suggested the particle nature of all electromagnetic
radiations which is popularly known as Quantum theory.
• Max Planck suggested that a body or atoms and molecules could emit or
absorb energy only in discrete quantity and not in the continuous form.
• He gave the name quantum to the smallest quantity of energy that can
be emitted or absorbed in the form of electromagnetic radiation. A
quantum of energy emitted or absorbed in the form of radiation is called
as photon.
• The energy (E) of a quantum or photon of radiation is proportional to its
frequency () and is expressed by equation
E α  or E = h
Where, E = Energy,  = frequency of radiation.
The proportionality constant “h” is known as Planck's constant and its value
is 6.626 X 10-34 J.s = 6.626 X 10-27 erg.s = 6.62 × 10–34 Kg m2s-1 = 4.135 × 10–15 eV.s.

11. • According to Max Planck, a body can absorb or emit only one
photon or some whole number multiple of it. Energy less than a
quantum or h can neither be absorbed nor be emitted.
• The energy can take any one of the values from the following set,
but can not take on any value between them.
E= h, 2h, 3h………nh i.e. E = nh
Thus, the energy is said to be quantized and n is called
as the quantum number.
Planck’s Quantum Theory

12. James Clarke Maxwell in 1873 revealed that light waves or electro
magnetic waves are associated with oscillating electric and
magnetic fields.
Characteristics of Electromagnetic radiation

13. Propagation of Electromagnetic radiation

14. Properties of electromagneticwaves
1. Light or EMR is a transverse wave and propagates through space in the
form of a wave as crest and trough. These waves are described to be
associated with oscillating electric and magnetic fields which are
perpendicular to each other and both are perpendicular to the
direction of the wave.
2. Electromagnetic waves do not require any specific medium for its
propagation and can travel through any matter or in vacuum.
3. In vacuum all types of electromagnetic radiations, regardless of wavelength
or frequency, travel at the same speed i.e 3.0 X 108m/s. This is called
speed of light and is given by symbol “c”.
4. There are many types of electromagnetic radiations, which differ from
one another in their wavelength or frequency. These constitute what is
called electromagnetic spectrum.
5. Electromagnetic spectrum is a continuous spectrum (i.e. the radiations
are diffused with each other) consisting of the following radiations.
Cosmic ray, Gamma ray ( – ray), X – ray, Ultraviolet ray, Visible ray,
Infrared ray, Microwave and radio wave. (C G X UV V IR M R)

15. Properties of electromagneticwaves
Increasing order of wavelength of these radiations is given as
Cosmic ray < Gamma ray < X – ray < Ultraviolet < Visible < Infrared <
Microwave < radio wave. (C G X UV V IR M R)
6. The energy transmitted from one body to another in the form of
radiations is called as radiant energy.
7. Electromagnetic radiations carry no charge.
8. An electromagnetic waves has no mass, but posses energy and
momentum (E = pc) and can exert some pressure called radiation
pressure. So, momentum can be given as, p = E/c = h/c = h/λ.
9. The energy of electromagnetic radiation is proportional to its
frequency or inversely proportional to its wavelength.
E α  or E = h = hc/λ = hc⊽ or E α 1/λ
Where, E = Energy,  = frequency of radiation, λ = wavelength of radiation and h
= Planck’s constant = 6.626 X 10-34 J.s = 6.626 X 10-27 erg.s
= 6.62 × 10–34 Kg m2s-1 = 4.135 × 10–15 eV.s.
All the electromagnetic radiations possesses certain wave characteristics such
as Wavelength (λ), F re q u e n c y ( ),Ve l o c i t y o r S p e e d ( v o r c ) ,
A m p l i t u d e ( A ) , I nte n s i t y ( I ) , Wave n u m b e r ( ⊽) a n d T i m e
Pe r i o d ( T ) .

16. Electromagnetic Spectrum

17. Properties of Electromagnetic Radiation
• Wavelength (λ):- It is the distance between two consecutive or successive
crests or trough of adjacent waves and is denoted as λ (Lambda).
Wavelength is expressed in the unit of Angstrom (Ao).
1Ao = 10-8 cm = 10-10m .
• Frequency ():- It is the number of cycles of the wave passing a given point per
second and is denoted as γ (nu), It is usually expressed in the unit of Hertz (Hz).
1Hertz (Hz) = 1 Cycle per Second or 1 Sec-1.
So, Frequency,  = c/λ = c⊽, where C = velocity of light & ⊽ = wave number.
• Speed or Velocity (V):- It is the distance travelled by a wave in one second and
is denoted as V. It is expressed in the unit of cm.sec-1 or m.sec-1.
• Amplitude (A):- It is the height of the crest or depth of the trough of a wave
and is denoted as A. It also determines the intensity of the wave and is expressed
in the unit of Angstrom (Ao) or cm or m.
• Wave number (⊽):- It is defined as the number of wavelengths per centimeter
and is the reciprocal of wavelength in centimeter. It is expressed in the unit of
cm-1 or m-1.
So, Wave number, ⊽ = 1/ λ = /c
• Time Period (T):- It is defined as the time taken for the light wave to complete
one cycle of wave and is the reciprocal of frequency. It is expressed in the unit
of Secor min.
So, Time period, T = 1/ = λ/c

19. The Wave Nature of the Electron
In 1924, Louis de-Broglie had arrived at his hypothesis with
the help of Planck’s quantum theory and Einstein’s equation
of mass-energy relationship.
He derived a relationship between the magnitude of the
wavelength associated with a particle of mass ‘m’ moving
with a velocity ‘c’. The wavelength of electron can be
described by the de Broglie relationship.
According to Quantum theory, E = h = hc/λ
According to Einstein’s theory, E = mc2
So, by combining, we can get, E = mc2 = hc/λ or λ = h/mc
Where, h = Planck’s constant = 6.626 x 10-34J.s = 6.626 x 10-34erg.s
m = mass of the particle & c = velocity of light.
Similarly, For a particle moving with velocity ‘v’, we can get,
λ = h/mv = h/p (p = momentum of particle = mv) i.e. λ  1/p.

20. Heisenberg's Uncertainty Principle
• Werner Heisenberg in 1927 developed the concept of the
Uncertainty Principle.
• “It is not possible to determine simultaneously both the exact
position and exact momentum of a microscopic particle such as
an electron with absolute accuracy or certainty”.
• Mathematically, it can be given as in equation
Where, ∆x = uncertainty in position, ∆p = uncertainty in momentum, ∆v= uncertainty in
velocity & ћ = h/2π = Reduced Planck’s Constant = 1.0546 × 10–34 J.s = 1.0546 × 10–27 erg.s
= 6.5821 × 10–16 eV.s.
• Importance of Heisenberg's Uncertainty Principle is only for moving micro
scopic object.

21. Photoelectric Effect
• In 1905, Albert Einstein explained the phenomenon of
photoelectric effect according to which, when a radiation of
suitable frequency or wavelength strikes the surface of certain
metals, electrons are ejected and such electrons ejected are
called as photoelectrons.

22. The observations of photoelectric effect were:
1. The electrons are ejected from the metal surface as soon as the
beam of light of a particular wavelength strikes the surface. i.e.
there is no time lag between the striking of light beam and the
ejection of electrons from the metal surface.
2. The number of electrons ejected is proportional to the intensity
or brightness of light.
3. For each metal, there is a characteristic minimum frequency o (
also known as threshold frequency) below which photoelectric
effect is not observed. At a frequency  > 0 , the ejected electrons
come out with certain kinetics energy.
4. Photons that have at least the minimum amount of energy called as
threshold energy (E0) having threshold frequency (0) will cause
the ejection of electrons from the metal surface.
Observations of Photoelectric Effect

23. The observations of photoelectric effect were:
5. The threshold energy (E0) required for ejection of electrons varies
from metal to metal.
6. Photoelectric current (no. of electrons emitted per second)
increases with the intensity of radiation but does not depend on
the frequency of radiation.
7. Kinetic energy of the ejected electrons is independent of the
intensity of radiation but depends on the frequency of radiation.
8. According to classical mechanics, an electron would be ejected as
soon as it could possess enough energy to overcome the force
which binds it to the metal. This energy is called as the work
function of the metal () and can be given as  = h0.
Observations of Photoelectric Effect

24. Mathematically, Photoelectric effect can be written as
Total energy of incident photon
= Binding energy + Kinetic energy of ejected electrons
Or Total energy of incident photon (h)
= Work function () + Kinetic energy (1/2mv2)
Or h =  + 1/2mv2
Or h = h0 + 1/2mv2
Or 1/2mv2 = h - h0
Or Kinetic energy (1/2mv2) = h( - 0)
Work function is the threshold energy just required to remove
the electron out of its nuclear attraction without giving it any
velocity.
Expression for Photoelectric Effect

25. Davison and Germer in the year 1927, deduced a relationship between
wavelength (λ) of an electron and the potential difference applied (V) on
it with the help of De-Broglie’s wave equation.
Consider an electron with mass (m) and charge (e) is accelerated through
a potential difference of V volts and it will start moving with a velocity
(v).
Thus, the kinetic energy of the electron = 1/2mv2
Electric force applied on the electron = e.V
But, the kinetic energy gained by the electron is due to the electric force
applied or potential difference applied on the electron.
i.e.
But, from De-Broglie’s equation,
Relationship between wavelength & Potential Difference

26. Concept of Black Body radiation
• An ideal body, which emits and absorbs radiation
of all frequency or wavelengths uniformly, is
called a black body and radiation emitted by
such a body is called black body radiation.
• In practice, no such ideal blackbody exists.
• Carbon black approximates fairly closely to the
black body.
• A good physical approximation to a black body is
a cavity with a tiny/pin hole, which has no other
opening at constant temperature. Any radiation
entering the hole will be reflected by the cavity
walls and will be eventually absorbed by the
walls.
• We can consider different planets and stars as
blackbodies.

27. • The amount of light emitted from the
black body and its spectral distribution
depends on its absolute temperature.
• At a particular temperature, intensity
of radiation emitted increases with
the increase in the wavelength, as
shown in figure.
• Also, as the temperature increases,
maxima of the curve shifts towards
the shorter wavelength.
Concept of Black Body radiation

28. Wien’s Displacement law
Wien's displacement law states that the wavelength corresponding
to maximum energy density (total energy per unit volume) in the
black body is inversely proportional to its absolute temperature.
Mathematically, Wien's displacement law can be given by:
λmax  1/T or λmax = b/T or λmax.T = b (constant)
where T is the absolute temperature.
b is a constant of proportionality called Wien's displacement constant,
equal to 0.29 × 10–2 mK or 2.9 × 10−3 m⋅K or b ≈ 2898 μm⋅K.
This is an inverse relationship between wavelength and temperature.
So, the higher the temperature, the shorter or smaller the wavelength of
the thermal radiation & the lower the temperature, the longer or larger
the wavelength of the thermal radiation.

29. Stefan – Boltzmann’s law
The Stefan - Boltzmann law describes the power radiated from a
black body in terms of its temperature.
Stefan – Boltzmann law states that the total energy radiated per
unit surface area of a black body across all wavelengths per unit
time (also known as emittance of the blackbody) is directly
proportional to the fourth power of the blackbody’s absolute
temperature, T.
Mathematically, Statement of Stefan – Boltzmann law can be given by:
E  T4 or E = σT4
Where T is the absolute temperature.
σ is a constant of proportionality called Stefan – Boltzmann’s constant
its value is 5.67 × 10–8 Wattm–2K–4 = 5.67 × 10–8 Joules–1m–2K–4

30. Rayleigh - Jeans law
This law was suggested by Lord Rayleigh and James Jeans in 1905 and is
based on the oscillating electromagnetic waves within the cavity of a
black body.
According to Rayleigh - Jeans law, the energy density (dE) in the
wavelength range of λ to (λ + δλ) is given by
Where f is energy per unit volume per unit wavelength.
λ is the wavelength of radiation, T is the absolute temperature.
k = Boltzmann’s Constant = 1.38 × 10–23 Kgm2s–2K–1.
Thus, the power emitted in this wave length range δλ is proportional to the
energy density.
But, this law was found to be valid for longer wavelengths and fails at shorter
wavelengths.

31. Modified Rayleigh - Jeans law (Planck)
To overcome the drawback of Rayleigh – Jeans law, Max Planck has
derived the equation for energy density in terms of quantization of
energy & can be given by
Where f is energy per unit volume per unit wavelength.
λ is the wavelength of radiation, T is the absolute temperature.
h = Planck’s constant = 6.626 X 10-34 J.s = 6.626 X 10-27 erg.s = 6.62 × 10–34 Kg m2s-1
k = Boltzmann’s Constant = 1.38 × 10–23 Kgm2s–2K–1.
This expression is valid for both longer and shorter wavelength range.

32. Proof of Planck’s Equation
For shorter wavelength, λ → 0 & is large &
i.e. λ → 0, ϒ → ∞ & f → 0
Hence,
So, energy density tends to zero at shorter wavelengths or high frequency.
Similarly, For longer wavelength,
This expression satisfies the Rayleigh – Jeans equation.

33. Characteristics of Visible light (VIBGYOR)
• The speed of light depends upon the nature of the medium
through which it passes. As a result, the beam of light is
deviated or refracted from its original path as it passes from one
medium to another medium. When light travels from one
medium to another, the wavelength changes but the frequency
does not change. As the light travel into the denser medium,
they slow down and wavelength decreases.

34. • A ray of white light is spread out in a series of colour bands
called spectrum.
• The light of red colour which has longest wavelength is deviated
the least, while the violet light which has shortest wavelength is
deviated the most.
• The spectrum of white light, that we can see ranges from violet at
7.50 X 1014 Hz to red at 4.0 X 1014 Hz. Such a spectrum is called a
continuous spectrum. Continuous because Violet merges into
indigo, Indigo merges into blue and so on.

35. Quantum Numbers
&
Shapes of Orbitals

36. “The QuantumNumbers”
Quantum numbers are the numbers which explains
the complete address of each electron in an atom. It
explained the arrangement and movement of
electrons, spectral lines of poly electronic atoms and
gave an acceptable model of an atom.

37. FourQuantumNumbers -An Electron’s Address
• Principal Quantum Number (n) - Specifies the main energy level
(Neil Bohr) (shell or orbit)
• Azimuthal Quantum Number (l) - Gives information about the
(Arnold Sommerfeld) sub energy level or sub shells or
sub orbits (orbital)
• Magnetic Quantum Number (m) - Spatial orientations of an
(Lande) orbital
• Spin Quantum Number (s) - Spin movement of electrons about its
(Uhlenbeck & Goudsmit) own axis

38. PrincipalQuantumNumber (n)
Size & Energy of an orbit/shell
and n = 1, 2, 3, 4,….
Greater value of n represents Bigger
orbits with higher energies.
Distance from the nucleus also
increases.

39. Principal QuantumNumber (n)
Total No. of Electrons in an orbit = 2n2
(According to Bohr – Bury Rule)
Value of n
Name of
Shell
Total No. of Electrons
2n2
n=1 K 2(1)2 = 2
n=2 L 2(2)2 = 8
n=3 M 2(3)2 = 18
n=4 N 2(4)2 = 32

40. AzimuthalQuantumNumber (l)
Each energy level is divided into sub levels.
“l” defines the shape of sub energy level or orbital.
For a given value of n, the energies of the sub shells
follow the order: s (l = 0) < p (l = 1) < d (l = 2) < f (l = 3) < g (l = 4)
No of Electrons in a sub shell is governed by 4l + 2 rule.
Name
Sub No. of
l
level electrons
0 s Sharp 2
1 p Principal 6
2 d Diffused 10
3 f Fundamental 14

41. Relationship between n & l
l = 0 → (n - 1)
Orbit n l Orbitals
No. of
electrons
K 1 0 1s 2
L 2 0, 1 2s, 2p 2+6 = 8
M 3 0, 1, 2 3s, 3p, 3d 2+6+10= 18
N 4 0, 1, 2, 3 4s, 4p, 4d, 4f 2+6+10+14=32

42. n=1 , K shell
n=2 , L shell
n=3 , M shell
n=4 , N shell
1s (2 electrons)
2s (2 electrons)
2p (6 electrons)
3s (2 electrons)
3d (10 electrons)
3p (6 electrons)
4s (2 electrons)
4d (10 electrons)
4p (6 electrons)
Increasing
Energy
&
Size
n l
4f (14 electrons)

43. MagneticQuantumNumber (m)
It explains the effect of external magnetic field on the
orientation of electrons in an orbital.
Sub shells split up into degenerate orbitals (having same
energy & size) in a magnetic field.
Each degenerate orbital can hold a maximum of 2 electrons.
Maximum number of orbitals that can accommodated in a
Shell/Orbit is governed by n2 rule where n is the principal
quantum number.
The value of Magnetic Quantum Number can be given as, m
= 2l + 1 and the value of m varies from – l to + l through 0,
where l is the azimuthal quantum number.

44. Relationship between l & m
m = -l  0+l
l
l = 0, s
l = 1, p
l = 2, d
Degenerate
orbitals
1
3
5
l = 3, f
m
0
-1, 0, +1
-2, -1, 0, +1,
+2
-3, -2, -1, 0,
+1, +2, +3
7
No. of
electrons
2
2+2+2 = 6
2+2+2+2+
2 = 10
2+2+2+2+
2+2+2 =
14

45. n=1 , K shell
n=2 , L shell
n=3 , M shell
n=4 , N shell
l =0, 1s
l =0, 2s
l =1, 2p
l =0, 3s
Increasing
Energy
&
Size
m=0
m=-1 m=0 m=+1
m=0
m=-1 m=0 m=+1
m=0
m=-2 m=-1 m=0 m=+1 m=+2
m=-1 m=0 m=+1
m=0
l =0, 4s
l =2, 3d
l =1, 3p
l =1, 4p
n l
l =3, 4f
l =2, 4d
m
m=-3 m=-2 m=-1 m=0 m=+1 m=+2 m=+3
m=-2 m=-1 m=0 m=+1m=+2

46. Sub level and Capacity of Each Main Energy Level
Main
Energy
Level
No. of
Sublevel
Identity of
Sublevels
No. of
Orbitals
(n2)
Max. No. of
Electrons
(2n2)
1 1 1s 1 2
2 2
2s 1 2
2p 3 6
3s 1 2
3 3 3p 3 6
3d 5 10
4s 1 2
4 4
4p
4d
3
5
6
10
4f 7 14

47. SpinQuantum Number (s)
Direction of spin of an electron.
Electron which rotates around the nucleus also
rotates around its own axis.
This is called as self rotation.
Electrons can have either Clockwise (50%) or anti
clockwise (50%) rotation.
s = + 1/2 (↑) for Clockwise rotation.
s = – 1/2 (↓) for Anticlockwise rotation.

48. Spin of electron
Associated with magnetic field

49. OrbitalDiagrams
We often represent an orbital as a square and the
electrons in that orbital as arrows.
The direction of the arrow represents the spin of
the electron.
Orbital with
1 electron
Unoccupied
orbital
Orbital with
2 electrons

50. Increasing
Energy
1s2
2p6
2s2
3d10
3p6
3s2
4f14
4d10
4p6
4s2
K shell
N shell
M shell
L shell

51. Orbit
The circular path of an
electron around the
nucleus is called an
orbit.
The orbit or shells are
denoted by K, L, M, N
etc

52. ElectronCloud
A cloud showing the
probability of finding the
electron in terms of charged
cloud around the nucleus is
called Electron Cloud.

53. Atomic Orbitals: s, p,d, f
Atomic orbitals are regions of space where the
probability of finding an electron about an atom is
highest.
s – orbital  spherical shape
p – orbital  dumb-bell shape
d – orbital  clover leaf shape or
double dumbbell
shape
f – orbital  double clover leaf or
too complex shape

54. s - orbital - spherically symmetrical
l = 0 & m = 0

55. Shapes of p – orbitals – Dumbbell
l = 1 & m = -1, 0, +1

56. Shapes of d – orbitals – Double Dumbbell
l = 2 & m = -2, -1, 0, +1, +2

57. Shapes of f – orbitals – Complex
l = 3, m = -3,-2,-1,0,+1,+2, +3

58. Distinction: Orbit & Orbital
Sl No ORBIT ORBITAL
1
It is a well defined circular path around the
nucleus in which electron revolves.
It is a region in three dimensional space
around the nucleus where the probability of
finding the electrons is maximum.
2
All the permitted orbits of an electron is
circular in shape.
Orbitals are different in their shape such as
s, p and d orbitals are spherical, dumbbell
and double dumbbell in shape.
3
It represents that an electron moves around
the nucleus in one plane.
It represents that an electron can move
around the nucleus along the 3 - dimensional
space (i.e. x, y and z axes).
4
Maximum number of electrons in an orbit is
2n2
where n is tne number of orbit.
Maximum number of electrons in an orbital
is 2 with opposite spins.
5
it represents that the position and the
momentum of an electron can be known
simultaneously with certainty which is against
the Heisenberg's uncertainty principle.
it represents that the position as well as the
momentum of an electron can not be known
simultaneously with certainty which is
accordance with Heisenberg's uncertainty
principle.
6
The shape of molecules can not be explained
by an orbit as they are non directional by
nature.
The shape of molecules can be found out by
orbitals as they are directional by nature
except s - orbitals.

59. Pauli’s Exclusion Principle
This rule was suggested by Wolfgang Pauli in 1952.
According to this principle, No two electrons in an atom can
have all the four Quantum numbers alike or same.
or
It can also be stated as only two electrons may exist in the
same orbital and these electrons must have opposite spins.
This principle is called the exclusion principle since it
excludes the possibility of accommodating more than 2n2
electrons in an orbit (where ‘n’ is the no. of orbit).

60. Hund’s Rule of Maximum Multiplicity
This rule was suggested by Friedrich Hund in 1925.
This rule deals filling of electrons in the orbitals of the
same sub shells of equal energy called degenerate
orbitals.
According to Hund’s rule, No electron pairing takes place
in degenerate p, d and f sub shells until each degenerate
orbital in the given sub shell contains one electron or
singly occupied.
Thus, exactly half filled and fully filled subshells are
most stable due to symmetry i.e. np3, np6, nd5, nd10, nf7
& nf14 configurations are most stable and elements try
to have stable configurations.

62. Auf-Bau’s Principle or Building up Principle
This rule was suggested by Neil Bohr & Wolfgang Pauli
in early 1920s. According to this rule, electron of an
atom first enters into the empty orbital of lowest
energy. When the lowest energy orbitals are
completely filled, electrons then enters into the orbital
of higher energy and this way electrons are arranged in
other orbitals based on their energy.
The arrangement of electrons in the orbitals/subshell of
an atom is governed by (n + l) rule, where n is the
principal quantum number & l is the azimuthal
quantum number.

63. Auf-Bau’s Principle or Building up Principle
Lower (n + l) values correspond to lower orbital energy and
shall be filled first than the orbital with higher (n + l) value.
Example:- 4s orbital has (n + l) rule = 4 and 3d orbital has (n
+ l) rule = 5. Hence, 4s orbital shall be filled prior to the
filling of 3d orbital.
But, if the (n + l) value is same for two different
orbitals/subshell, then the orbital with lower value of ‘n’ is
assumed to have lower energy and shall be filled up first.
Example:- 4s orbital has (n + l) rule = 4 and 3p orbital has (n
+ l) rule = 4. Hence, 3p orbital shall be filled prior to the
filling of 4s orbital due to lower value of ‘n’.

69. ELECTRONIC CONFIGURATION OF 30 ELEMENTS
ELEMENT NAME ATOMIC NUMBER SYMBOL ELECTRON ARRANGEMENT ELECTRONIC CONFIGURATION
Hydrogen 1 H 1 1s1
Helium 2 He 2 1s2
Lithium 3 Li 2, 1 1s2
2s1
Beryllium 4 Be 2, 2 1s2
2s2
Boron 5 B 2, 3 1s2
2s2
2p1
Carbon 6 C 2, 4 1s2
2s2
2p2
Nitrogen 7 N 2, 5 1s2
2s2
2p3
Oxygen 8 O 2, 6 1s2
2s2
2p4
Fluorine 9 F 2, 7 1s2
2s2
2p5
Neon 10 Ne 2, 8 1s2
2s2
2p6
Sodium 11 Na 2, 8, 1 1s2
2s2
2p6
3s1
Magnesium 12 Mg 2, 8, 2 1s2
2s2
2p6
3s2
Aluminum 13 Al 2, 8, 3 1s2
2s2
2p6
3s2
3p1
Silicon 14 Si 2, 8, 4 1s2
2s2
2p6
3s2
3p2
Phosphorus 15 P 2, 8, 5 1s2
2s2
2p6
3s2
3p3
Sulphur 16 S 2, 8, 6 1s2
2s2
2p6
3s2
3p4
Chlorine 17 Cl 2, 8, 7 1s2
2s2
2p6
3s2
3p5
Argon 18 Ar 2, 8, 8 1s2
2s2
2p6
3s2
3p6
Potassium 19 K 2, 8, 8, 1 1s2
2s2
2p6
3s2
3p6
4s1
Calcium 20 Ca 2, 8, 8, 2 1s2
2s2
2p6
3s2
3p6
4s2
Scandium 21 Sc 2, 8, 9, 2 1s2
2s2
2p6
3s2
3p6
3d1
4s2
Titanium 22 Ti 2, 8, 10, 2 1s2
2s2
2p6
3s2
3p6
3d2
4s2
Vanadium 23 V 2, 8, 11, 2 1s2
2s2
2p6
3s2
3p6
3d3
4s2
Chromium 24 Cr 2, 8, 13, 1 1s2
2s2
2p6
3s2
3p6
3d5
4s1
Manganese 25 Mn 2, 8, 13, 2 1s2
2s2
2p6
3s2
3p6
3d5
4s2
Iron 26 Fe 2, 8, 14, 2 1s2
2s2
2p6
3s2
3p6
3d6
4s2
Cobalt 27 Co 2, 8, 15, 2 1s2
2s2
2p6
3s2
3p6
3d7
4s2
Nickel 28 Ni 2, 8, 16, 2 1s2
2s2
2p6
3s2
3p6
3d8
4s2
Copper 29 Cu 2, 8, 18, 1 1s2
2s2
2p6
3s2
3p6
3d10
4s1
Zinc 30 Zn 2, 8, 18, 2 1s2
2s2
2p6
3s2
3p6
3d10
4s2