These notes covers chemistry chapter 2nd of class 11th which are strictly according to CBSE & state board syllabus.The contents covered are Model of atom, electronic configuration & many more..

1. Page 1
UNIT 2
STRUCTURE OF ATOM
DISCOVERY OF ELECTRON, PROTON & NEUTRON
Matter is made up of tiny particles called atom. Atoms are further made of three fundamental
particles or sub-atomic particles called electron, proton and neutron.
Proton E.Goldstein
ईट पग नाचे
Electron James Chadwick
J.J. Thomson Neutron
ELECTRON (e-) PROTON (p) NEUTRON (n)
Discovery J.J. Thomson E. Goldstein James Chadwick
Nature of Charge Negative Positive Neutral
Amount of Charge −1.6 × 1019
𝐶 +1.6 × 1019
𝐶 zero
Mass 9.1 × 10−31
𝑔 1.67 × 10−27
𝑔 1.67 × 10−27
𝑔
 Electrons are the subatomic particles which carry unit negative charge. Electrons were
discovered using Cathode ray Discharge tube.
 Cathode rays are the stream of negatively charged particles (electrons) which start from
Cathode & end at anode.
 Protons are the subatomic particles which carry unit positive charge. Protons were
discovered using Modified Cathode ray Discharge tube.
 Anode rays are the stream of positively charged particles (protons). Anode rays are not
originated from anode; they are originated between the space of anode & cathode. Anode
rays are also called Canal Ray as it passes through the holes or canals in the cathode.
Add to your knowledge: Cathode Ray Tube (CRT Screen) is used as Computer & TV Monitor.
 Neutrons are the subatomic particle carries no charge.
ATOMIC NUMBER & MASS NUMBER
Atomic Number (Z) of an element is equal to the number of protons present in the atom.
Mass Number (A) of an element is equal to the total number of protons & neutrons present in
the atom OR number of nucleons present in the atom (as Protons & Neutrons are collectively
called nucleons).
𝑋𝑍
𝐴
Remember: Number of Protons = 𝒁
Number of Neutrons= 𝑨 − 𝒁
Number of Electrons = 𝒁 (𝑓𝑜𝑟 𝑛𝑒𝑢𝑡𝑟𝑎𝑙 𝑎𝑡𝑜𝑚 𝑿)
= 𝒁 − 𝒏 (𝑓𝑜𝑟 𝑐𝑎𝑡𝑖𝑜𝑛 𝑿 𝒏+
)
= 𝒁 + 𝒏 (𝑓𝑜𝑟 𝑎𝑛𝑖𝑜𝑛 𝑿 𝒏−
)

2. Page 2
THOMSON’S MODEL OF ATOM
This model proposed that atom is considered as a uniformly positively charged sphere &
electrons are embedded in it. This model is also called as plum pudding, raisin pudding &
Watermelon Model.
Important features:
 Mass of the atom is considered to be distributed evenly over the atom.
 Atom is overall neutral.
Drawback:
 This model failed to explain the observations of Rutherford’s α-particle scattering
experiment.
Remember: α-particles are the nucleus of He-atom i.e. 𝑯𝒆 𝟐+
.
RUTHERFORDS’S MODEL OF ATOM
This model proposed that atom has two parts Nuclear & Extra-nuclear part; in nuclear part there
is nucleus which comprises of protons & neutrons and in extra-nuclear part, electrons are
revolving in the circular orbit.
Important features:
 Atom posses a highly dense, positively charge sphere called nucleus.
 Entire mass of atom is concentrated inside the nucleus.
 Electrons & nucleus are held together by electrostatic force of attraction.
Drawbacks:
 As electron revolves around the nucleus in circular orbit and circular motion is an
example of uniformly accelerated motion. So, there must be acceleration in moving
electron around the nucleus & as we know an accelerated charge particle emit
Electromagnetic radiation. Thus, the orbit should shrink but this doesn’t happen.
 It doesn’t say anything about electron distribution around the nucleus.
 It couldn’t explain the stability of atom.
Remember: Size of nucleus ≈ 10−15
𝑚 (1 𝑓𝑒𝑟𝑚𝑖)
Size of atom ≈ 10−10
𝑚 (1 𝐴ᵒ 𝑜𝑟 1 𝐴𝑛𝑔𝑜𝑠𝑡𝑟𝑜𝑚)
ISOTOPES, ISOBARS & ISOTONES
 Isotopes of an element are the atoms of the element with same atomic number but
different mass number e.g. 𝐻, 𝐻 & 𝐻1
3
1
2
1
1
.
 Isobars are the atoms of the different elements with same mass number but different
atomic number e.g. 𝐶6
14
& 𝑁.7
14
 Isotones are the atoms of the different elements containing same number of neutrons e.g.
𝐶 & 𝑁.7
14
6
13
Remember: Last Third alphabet
ISOTOPES Number of Protons are same
ISOBARS Atomic Mass same
ISOTONES Number of Neutrons are same.

3. Page 3
ISOELECTRONIC SPECIES
Isoelectronic Species are the species of different elements containing same number of electrons
e.g. 𝑂2−
, 𝐹−
, 𝑁𝑒, 𝑁𝑎+
& 𝑀𝑔2+
𝑎𝑙𝑙 ℎ𝑎𝑣𝑒 10 𝑒𝑙𝑒𝑐𝑡𝑟𝑜𝑛𝑠.
ELECTROMAGNETIC RADIATION
These Radiations are associated with oscillating Electric & Magnetic field; these two fields are
perpendicular to each other & also to the direction of propagation. These Radiations travel with
the speed of light & do not require any material medium for their propagation.
Characteristics of electromagnetic Radiation
1. Wavelength ( ) is defined as the distance between two consecutive crest & trough.
2. Frequency ( ) is defined as the time taken to complete one oscillation.
3. Velocity (𝒄) is defined as the distance travelled by the wave in one second.
Relation between wavelength, frequency & velocity 𝒄 =  
4. Wave number ( ) is defined as the reciprocal of wavelength.
5. Amplitude (A) is defined as the height of crest & depth of trough.
ELECTROMAGNETIC SPECTRUM
Arrangement of all the electromagnetic radiations in the increasing order of their wavelengths or
decreasing order of their frequencies is called electromagnetic spectrum.
गा ज़रा उल्टा-वुल्टा in my Radio
Gamma rays X-rays Utravoilet Visible Infrared Microwave Radio wave
Add to your knowledge: Visible light has wavelength range of 400 nm to 750 nm.
PARTICLE NATURE OF ELECTROMAGNETIC RADIATION
Planck’s quantum theory states that the energy emitted or absorbed in the form of small
packets of energy called ‘quanta’. In case of light, the quantum of energy is often called photon.
This energy is directly proportional to the frequency of radiation.
𝐸 ∝ 𝜈
𝐸 = ℎ𝜈 𝑤ℎ𝑒𝑟𝑒 ℎ 𝑖𝑠 𝑐𝑎𝑙𝑙𝑒𝑑 𝑝𝑙𝑎𝑛𝑐𝑘′
𝑠 𝑐𝑜𝑛𝑠𝑡𝑎𝑛𝑡 = 6.6 × 10−34
𝑗𝑜𝑢𝑙𝑒 − 𝑠𝑒𝑐
OR 𝐸 = ℎ
𝑐
𝜆
The energy is always absorbed or emitted as the integral multiple of this quantum.
Photo Electric Effect is the phenomenon of ejection of electron from the metal surface when a
radiation of suitable frequency falls on it.
Experimental Results
 Electrons are ejected immediately as light falls on the metal surface.
 Number of electrons ejected is directly proportional to intensity or brightness of light.
 If the frequency of incident radiation is less than some critical value, no photo electron
will be ejected.
Increasing wavelength
Decreasing frequency

4. Page 4
Threshold frequency (𝜈o) is the characteristics minimum frequency below which photo
electric effect is not observed.
Work Function (𝝋 𝒐) is the minimum energy required to eject electrons.𝝋 𝒐 = 𝒉𝝂 𝐨
Einstein Explanation: The energy of incident photon is used to eject electron from the
metal. If the photon has energy more than the energy required to eject the electron, the excess
energy is taken up by the electron as its kinetic energy.
𝑬 = 𝝋 𝒐 + 𝑲. 𝑬.
𝒉𝝂 = 𝒉𝝂 𝐨 +
𝟏
𝟐
𝒎𝒗 𝟐
Black Body is an ideal body, which emits & absorbs all frequencies & the radiations emitted by
such body is called Black Body Radiation.
ATOMIC SPECTRA
Ordinary white light consist of waves with all wavelengths in visible range, a ray of white light is
spread out into a series of coloured band called Spectrum.
 The spectrum consists of all the wavelength is known as Continuous spectrum.
 The spectrum in which only specific wavelength is present is known as Line spectrum.
 The spectrum of radiation emitted by a substance that has absorbed energy is called an
Emission spectrum.
 The spectrum obtained when radiation is passed through the sample of a material, the
sample absorbs radiation of certain wavelengths is known as Absorption spectrum.
Add to your knowledge: Line emission spectrum is used to identify unknown atoms.
The element helium (He) was discovered in the sun by spectroscopic method.
LINE SPECTRA FOR HYDROGEN
 Line spectrum for hydrogen is helpful in the development of atomic structure.
 Hydrogen spectrum is observed when electric discharge is passed through gaseous
Hydrogen & Hydrogen atom get excited.
Series n1 n2 Spectral region
Lyman 1 2, 3… Ultraviolet
Balmer 2 3, 4… Visible
Paschen 3 4, 5… Infrared
Brackett 4 5, 6… Infrared
Pfund 5 6, 7… Infrared
Rydberg’s Formula
𝝂̅ = 𝟏𝟎𝟗, 𝟔𝟕𝟕(
𝟏
𝒏 𝟏
𝟐 −
𝟏
𝒏 𝟐
𝟐 ) 𝒄𝒎−𝟏
[109,677 cm-1 is Rydberg Constant for Hydrogen]
Remember:
1. Number of spectral line in the spectrum when electron comes from nth level to ground
level.
𝒏𝒐. 𝒐𝒇 𝒔𝒑𝒆𝒄𝒕𝒓𝒂𝒍 𝒍𝒊𝒏𝒆𝒔 =
𝒏( 𝒏 − 𝟏)
𝟐
2. To find longest wavelength (First Line of Series) & shortest wavelength (Last line of
Series) transition for any series- n1 is fixed and n2 is minimum & maximum respectively
e.g. longest wavelength transition for Balmer series n1 = 2 (Fixed) & n2 = 3 (Minimum)
and for shortest wavelength transition n1 = 2 (Fixed) & n2 = ∞ (Maximum).

5. Page 5
BOHR’S MODEL OF ATOM
 Applicable only for one electron system like H, He+ & Li2+.
 The electron in the Hydrogen atom revolves around the nucleus in circular orbits of fixed
radius & energy called energy shells or energy levels.
 The energy of an electron in a particular orbit remains constant. These orbits are called
stationary states.
 Energy is emitted or absorbed when an electron jumps from higher energy level to lower
energy level or vice-versa.
∆𝐸 = 𝐸2 − 𝐸1 = ℎ𝜈
𝜈 =
𝐸2 − 𝐸1
ℎ
 An electron can move only on those orbits for which its angular momentum is an
integral multiple h/2π.
𝑚𝑣𝑟 = 𝑛
ℎ
2π
Remember: The most stable state of an atom is called its ground state or normal state.
Any state of an atom that has higher energy than the ground state is called its
excited state.
MATHEMATICAL EXPRESSIONS
1. Energy of an electron in nth orbit (𝐸 𝑛 )
𝑬 𝒏 = −
𝟐. 𝟏𝟖 × 𝟏𝟎−𝟏𝟖
𝒁 𝟐
𝒏 𝟐
𝑱/𝒂𝒕𝒐𝒎 [
𝑍 = 𝑎𝑡. 𝑛𝑢𝑚𝑏𝑒𝑟
𝑛 = 𝑛𝑜. 𝑜𝑓 𝑜𝑟𝑏𝑖𝑡
]
= −
𝟏𝟑. 𝟔 × 𝒁 𝟐
𝒏 𝟐
𝒆𝑽/𝒂𝒕𝒐𝒎
= −
1312 × 𝑍2
𝑛2
𝑘𝐽/𝑚𝑜𝑙𝑒
2. Radius of an orbit (𝑟𝑛)
𝒓 𝒏 =
𝟎. 𝟓𝟐𝟗𝒏 𝟐
𝒁
𝑨°
3. Velocities of electrons moving in different orbits (𝑣 𝑛)
𝒗 𝒏 =
𝟐. 𝟏𝟖 × 𝟏𝟎 𝟔
× 𝒁
𝒏
𝒎/𝒔
LIMITATIONS OF THE BOHR’S MODEL
(i) It could not explain the spectra of atom containing more than one electron.
(ii) It fails to explain splitting of spectral line in the presence of electric field (Stark effect)
or magnetic field (Zeeman effect).
(iii)It ignored dual character of electron.
(iv)It contradicts Heisenberg’s uncertainty principle.
(v) It could not explain ability of atom to form chemical bonds and the geometry & shapes of
the molecule.

6. Page 6
DUAL BEHAVIOUR OF MATTER
Matter exhibits dual behavior Particle nature as well as wave nature.
de-BROGLIE HYPOTHESIS
Every object in motion has a wave character. The wave associated with material particles are
called matter waves or de-Broglie waves.
𝑊𝑎𝑣𝑒 𝑛𝑎𝑡𝑢𝑟𝑒 ∝
1
𝑃𝑎𝑟𝑡𝑖𝑐𝑙𝑒 𝑁𝑎𝑡𝑢𝑟𝑒
𝑤𝑎𝑣𝑒𝑙𝑒𝑛𝑔𝑡ℎ ∝
1
𝑚𝑜𝑚𝑒𝑛𝑡𝑢𝑚
𝜆 ∝
1
𝑝
𝝀 =
𝒉
𝒎𝒗
[ 𝑤ℎ𝑒𝑟𝑒 ℎ = 𝑝𝑙𝑎𝑛𝑐𝑘′
𝑠 𝑐𝑜𝑛𝑠𝑡𝑎𝑛𝑡]
𝝀 =
𝒉
√ 𝟐𝒎𝑬
[ 𝐸 = 𝐾𝑖𝑛𝑒𝑡𝑖𝑐 𝑒𝑛𝑒𝑟𝑔𝑦 𝑜𝑓 𝑡ℎ𝑒 𝑒𝑙𝑒𝑐𝑡𝑟𝑜𝑛]
Remember: The wavelength associated with ordinary object is so short that their wave
properties cannot be detected & the wavelengths associated with sub-atomic particles can
however be detected experimentally.
HEISENBERG’S UNCERTAINTY PRINCIPLE
It is impossible to measure simultaneously both the position & velocity of a microscopic particle
with absolute accuracy or certainly. (To observe an electron, we can illuminate it with light. This
light has enough energy to change the velocity of the electron).
∆𝒙 × ∆𝒑 ≥
𝒉
𝟒𝝅
[
∆𝑥 = 𝑢𝑛𝑐𝑒𝑟𝑡𝑎𝑖𝑛𝑖𝑛𝑡𝑦 𝑖𝑛 𝑝𝑜𝑠𝑖𝑡𝑖𝑜𝑛
∆𝑝 = 𝑢𝑛𝑐𝑒𝑟𝑡𝑎𝑖𝑛𝑖𝑛𝑡𝑦 𝑖𝑛 𝑚𝑜𝑚𝑒𝑛𝑡𝑢𝑚
]
∆𝒙 × ∆𝒗 ≥
𝒉
𝟒𝝅𝒎
[∆𝑣 = 𝑢𝑛𝑐𝑒𝑟𝑡𝑎𝑖𝑛𝑖𝑡𝑦 𝑖𝑛 𝑣𝑒𝑙𝑜𝑐𝑖𝑡𝑦]
If uncertainty in position (∆𝑥) is less, then uncertainty in momentum (∆𝑝) be large & vice-versa.
Heisenberg’s uncertainty principle rules out the existence of definite path or trajectories of
electrons. It therefore, means that the precise statements of position & momentum of electrons
have to be replaced by the statement probability that the electron has at given position &
momentum.
QUANTUM MECHANICAL MODEL OF ATOM
This model is based on a fundamental equation which is called Schrodinger wave equation.
Hˆ  = E [where Hˆ is a mathematical operator]
The Solution of Schrodinger wave equation gives the value of E & . The corresponding value
of  are called wave functions.
 gives us the amplitude of the wave & the value of  has no physical significance.
2 gives us the region in which the probability of finding the electron is maximum & it is called
probability density. This region around the nucleus is called orbital.
An orbital is defined as the three dimensional space around the nucleus within which the
probability of finding the electron of given energy is maximum.

7. Page 7
Orbit Orbital
1. Well defined circular path around the
nucleus in which the electron revolve.
1. Space around the nucleus within which the
probability of finding the electron of given
energy is maximum
2. Planar motion of an electron 2. Three dimensional motion of an electron.
3. Maximum number of electron in any orbit is
given by 2n2.
3. Maximum number of electron in any orbital
is given by 2.
4. Orbits are circular in shape. 4. Orbitals have different shapes.
QUANTUM NUMBERS
Each electron in an atom is identified in the terms of four quantum numbers:
1. Principal quantum number (n)
2. Azimuthal quantum number (l)
3. Magnetic quantum number (m)
4. Spin quantum number (s)
Principal quantum number (n) identifies shells; determine size & energy of the orbital.
n 1 2 3 4
Shell K L M N
Total no. of orbital 1 4 9 16
Maximum no. of electrons 2 8 18 32
Remember:
1) Total no. of orbital = n2
2) Maximum no. of electron in any shell = 2n2
Azimuthal quantum number (l) identifies sub-shell; determine shape, energy & angular
momentum of the orbital.
l s p d f g
Designation of sub-shell 0 1 2 3 4
Number of orbitals 1 3 5 7 9
Remember:
1) For a given value of n, l ranging from 0 to n-1.
n l Sub-shell notations
1 0 1s
2
0 2s
1 2p
3
0 3s
1 3p
2 3d
4
0 4s
1 4p
2 4d
3 4f
2) For a given value of l, Orbital angular momentum =√𝒍( 𝒍+ 𝟏)
𝒉
𝟐𝝅

8. Page 8
Magnetic quantum number (m) gives information about spatial orientation of the orbital with
respect to standard set of co-ordinate axis.
l 0 1 2
m 0 -1 0 1 -2 -1 0 1 2
Orbital s py pz px dxy dyz dz
2 dxz dx
2-y
2
Remember:
1) For a given value of l, m ranging from –l to +l.
2) No. of orbital in a sub-shell m = 2l+1.
Spin quantum number (s) gives information about orientation of the spin of the electron
present in any orbital.
Remember:
For a given value of m, s can have only two values i.e. +1/2 and -1/2. +1/2 identifies the
clockwise spin & -1/2 identifies the anticlockwise spin.
Add to your knowledge: s, p, d & f orbitals are derived from the certain series of alkali
metal spectroscopic lines as sharp, principal, diffuse, and fundamental.
Maximum no. of electrons that can be accommodated in a shell, its sub-shell and orbitals
Shell n
l
(0 to n-1)
m
(-l to +l)
s
(±𝟏/𝟐)
Electron
present
Total no. of
electrons
K 1 0 (1s) 0 (1s) ±1/2 2 2
L 2
0 (2s) 0 (2s) ±1/2 2
8
1 (2p)
-1 (2py) ±1/2
60 (2pz) ±1/2
+1 (2px) ±1/2
M 3
0 (3s) 0 (3s) ±1/2 2
18
1 (3p)
-1 (3py) ±1/2
60 (3pz) ±1/2
+1 (3px) ±1/2
2 (3d)
-2 (3dxy) ±1/2
10
-1 (3dyz) ±1/2
0 (3dz
2) ±1/2
+1 (3dxz) ±1/2
+2 (3dx
2
-y
2) ±1/2
SHAPES OF ORBITAL: s, p, d Shapes
Shape of s-orbital
For s-orbital l = 0 & m = 0. So, there is only 1 possible orientation. Therefore, s-orbitals
have spherical shape.
1s 2s

9. Page 9
Remember:
1) The region where the probability of finding the electron is zero is called nodal surface or
node.
2) In any ns orbital there are (n-1) nodes.
Shape of p-orbital
For p-orbital l = 1 & m = -1, 0 & +1. So, there are 3 possible orientations denoted as px, py & pz.
p-orbitals have dumb-bell shape.
y
x
z
px py pz
Shape of d-orbital
For d-orbital l = 2 & m = -2, -1, 0 & +1, +2. So, there are 5 possible orientations denoted as dxy,
dyz, dxz, dx
2-y
2& dz
2. d-orbitals have double dumb-bell shape.
y z z
x y x
dxy dyz dxz
y z
x
dx
2-y
2 dz
2
FILLING OF ORBITALS IN ATOMS
The distribution of electrons in various orbitals is known as electronic configuration.
𝒏𝒍 𝒙
Aufbau Principle states that in the ground state of the atoms, the orbitals are filled in order of
increasing energies.
n+l Rule orbitals with lower value of (n+l) value lower energy. If the two orbitals have the same
value of (n+l) then the orbital with the lower value of n will have lower energy.
The order in which the orbitals are filled is as follows:
1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p…

10. Page 10
1s
2s 2p
3s 3p
4s 3d 4p
5s 4d 5p
6s 4d, 5f 6p
7s 5d, 6f 7p
Hund’s Rule of maximum multiplicity states that electron pairing in p, d and f orbitals cannot
occur until each orbital of a given sub-shell contains ne electron each or is singly occupied.
 2px
22py
1
 2px
12py
12pz
1
Pauli’s Exclusion Principle states that an orbital can have two electrons & must have opposite
spins.
METHODS OF WRITING ELECTRONICS CONFIGURATIONS
Electronic configuration of Fluorine
1s22s22p5 2, 7
(Orbital method) (Shell method)
1s 2s 2p
(Box Method)
Stability of completely filled & half filled sub shell
(i) Symmetrical distribution of electron- the completely filled & half filled sub-shell have
symmetrical distribution of electrons.
(ii) Exchange energy- The two or more electrons with the same spin present in degenerate
orbitals (orbitals of same energy) of a sub-shell can exchange their position & energy
released due to exchange is called exchange energy.
The no. of exchanges is maximum when the sub-shell is either half filled or completely
filled.
Electronic configuration of Chromium
 [Ar]4s23d4
 [Ar]4s13d5
Electronic configuration of Copper
 [Ar]4s23d9
 [Ar]4s13d10
(Half filled stable) (Full filled stable)