Kerala Technical Board Polytechnic Colleges- Engineering Chemistry-II Semester- II Module-1:- Atomic structure-II and chemical bonding

1. By MUHAMMED SALMAN FARISH TP
Email: tpmsalman@gmail.com

3. Bohr model of atom

4. Postulates of Bohr atomic model
1. Electrons are revolving around the nucleus in a circular path
of fixed energy called orbits.
2. The angular momentum of an electron is an integral multiple
of
ℎ
2π
3. The energy of an electron in the orbit does not change with
time.
4. The frequency of radiation absorbed or emitted when
transition occurs between two stationary states is given by:
ν =
ΔE
ℎ
=
E2 – E1
ℎ

5. Demerits of Bhor atomic model
1) He explained the line spectra of hydrogen atom and
hydrogen like ions only.
2) He couldn’t explain the spectra of atoms other than
hydrogen.
3) He couldn’t explain Stark effect ( splitting of spectral
line in electric field ) and Zeeman effect (splitting of
spectral line in magnetic field).
4) Didn’t explain about the bonding of atoms and
formation of molecules.
5) He couldn’t explain the wave nature of atoms.

6. Dual nature of matter

7. de Broglie relation
 He stated, matter can exhibit dual behaviour i.e., both
particle and wave nature. He gave;
λ =
𝒉
𝒎𝒗
=
𝒉
𝒑
where, λ = wavelength
h = planks constant
m = mass
v = velocity
p = momentum

8. Heisenberg’suncertainityprinciple
 “it is impossible to determine the exact position and
exact momentum (or velocity) of a moving microscopic
particle like electron simultaneously”
Δx. Δp ≥
ℎ
4π
Δx. Δp ≥
ℎ
4πm
Where Δx =change in position
Δp = change in momentum

9. Orbit
• Circular path
around the
nucleus
• K,L,M,N,O
• 2-Diamensional
Orbital
• Area around the
nucleus where
probability to
find e- s are
maximum
• s,p,d,f
• 3-diamensinal

10. Quantum numbers
Principle quantum number
Azimuthal quantum number
Magnetic quantum number
Spin quantum number

11. Principle quantum numbers (n)
1. It gives the size the orbit.
2. It gives the energy of electron in an orbit.
3. It gives the shell in which the electron is found
for K shell, n = 1
for L shell, n = 2
for M shell, n = 3
for N shell, n = 4

12. Azimuthal quantum number (l)
1. It gives the shape of the orbital.
2. It gives the sub shell or sub level in which the electron is
located.
3. It also gives the orbital angular momentum of the
electron.
l can be from 0 to n-1
ie for n = 1, l = 0
for n = 2, l = 0, 1
for n = 3, l = 0, 1, 2

13. Magnetic quantum number (m)
 It gives information about the orientation of orbitals in
space. There are 2 l +1 possible values for m
for l = 0, m = -1,
for l = 1, m = -1,0,1
for l = 2, m = -2, -1, 0, 1, 2
for l = 4.... m = …………
It’s value from –l to + l

14. Spin quantum number (s)
 It is the only experimental Quantum number and it
gives the spin orientation of electrons. This spin may
be either clockwise or anticlockwise. So the values for
s may be +½ or -½. +½ represents clock-wise spin
and-½ represents anticlockwise spin.
 For all, S = +½ or -½

15. Shapes of orbitals
P orbital-
Dumb bell shape
S orbital-
Spherical

16. Chemical bonding
Ionic bond
• Eg: NaCl , KCl
Covalent bond
• Eg: F2,Cl2, O2, N2
Coordinate bond
• Eg: NH4
+, H3O+
Hydrogen bonding
• H2O , HF

17.  Ionic bond formed by the transfer of electrons (donate or
accept) between atoms and there by completing the octet.
It is mainly formed between electropositive and
electronegative elements. The constituent atoms are held
together by electrostatic force of attraction by the
formation of +ve (cation) and –ve (anions) ions.
 When an electron donated by an atom, get a +ve charge
 When an electron accepted by an atom, get a –ve charge
Ionic bond ( electrovalent bond)

18. Ionic bond ( electrovalent bond)
Cl
●●
●●
●●
●
Na ●
Na ●
 Eg;
11Na: 2, 8, 1 17Cl: 2, 8, 7 2,8 & 2,8,8
Na donated one electron to Cl and forms Na+ and Cl- ions,
and thereby forms NaCl molecule as a result of ionic bond.
Na + Cl → [Na]+[Cl]- ie, NaCl molecule.

19. Covalent bond
Single bond
• Sharing of 2 electrons(1 pair)
Double bond
• Sharing of 4 electrons(2 pair)
Triple bond.
• Sharing of 6 electrons (3 pair)

20. Single bond
 Bond formed by the sharing of 2 electrons between two
atoms
 eg: Cl + Cl → Cl2 or F+ F → F2
Consider the formation of Cl2,
17Cl :2, 8, 7
Each chlorine contains 7 electrons in last shell, so the two
atoms will share one electron each, ie, one pair of electrons
between them and forms Cl2 molecule with a single bond.

21. Formation of chlorine molecule.
●
Cl
●●
●●
●●
Cl
●●
●
●●
●●
+ Cl
●●
●
●●
Cl
●●
●●
●●
●
●●
or Cl
●●
●●
Cl
●●
●● ●●
●●

22. Double bond
 Sharing 4 electrons between atoms.
 Eg; O + O → O2. atomic no. of oxygen 8. so Electronic
Configuration will be 2, 6.
O
●●
●●
●●
O
●●
●●
●●
+ O
●●
●●
●●
O
●●
●●
●●
O
●●
●●
O
●●
●●
═or

23. Triple bond
N
●●●
●●
N
●●●
●●
+ N
●●●
●●
N
●●●
●●
or N
●●
N
●●
═─
Sharing 6 electrons between atoms.
Eg; N + N → N2. atomic no. of nitrogen 7.
Electronic Configuration 2, 5,

24. Coordinate bond
 A bond formed by donating electron pair by one atom but
shared between both atoms to complete the octet.
 eg; the formation of NH3→ BF3
N
│
│
│
H
H
H
●●
+ B
│
│
│
F
F
F N B
│
│
│
F
F
F
│
│
│
H
H
H →→

25. H -bond
 It is the attractive force which binds hydrogen atom and
highly electronegative atoms like F, O, N
1. Intermolecular H bond.
H bond present between two similar/different
molecules.eg,
H-F -----H-F-----H-F------H-F
2. Intra molecular H bond
H bond present within the same molecule.

26. Aufbau principle
 It states that the orbitals are filled in the increasing order
of their energies, in other words, electrons first occupy
the lowest energy orbital and then to higher energy
orbitals.
 The increasing energy order of orbitals as follows;
1s2, 2s2, 2p6, 3s2, 3p6, 4s2, 3d10, 4p6,…..

27. Pauli’s exclusion principle
 It states that no two electrons in an atom can have the
same set of four quantum numbers. i.e. an orbital can
accommodate a maximum of only 2 electrons with
opposite spin.
 If 2 electrons have same values for n, Ɩ and m, they should
have different values for s. i.e. if s = +
𝟏
𝟐
for the first
electron, it should be -
𝟏
𝟐
for the second electron.
ie, ↑↓
2s
↑↓ ↑↓ ↑↓
2p ↑↑ ↑↑ ↑↑
2p

28. Hund’sruleofmaximummultiplicity
 It states that electron pairing takes place only after
partially filling all the degenerate orbitals. Orbitals
having same energies are called degenerate orbitals.
 For example the electronic configuration of N is 1s2 2s2
2px
1py
1pz
1 and not 1s2 2s2 2px
2py
1
↑ ↑ ↑↑↓↑↓
2s1s 2px2, py1,
↑↓ ↑not ↑↓↑↓
1s 2s2px1, 2py1,2pz1