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11. ⮚ It is defined as the amount of energy required to break one mole of
bonds between two atoms in a gaseous state.
⮚ The unit of bond enthalpy is kJ mol–1. For example, the H – H bond
enthalpy in hydrogen molecule is 435.8 kJ mol–1.
H2(g) H (g) H (g)
+ : Δ H = 435.8 kJ/mol
Bond Enthalpy
12. ⮚ In case of polyatomic molecules, the measurement of bond strength is
more complicated.
⮚ For example : In case of H2O molecule, the enthalpy needed to break
the two O–H bonds is not the same.
⮚ Therefore in polyatomic molecules the term average bond enthalpy is
used.
Bond Enthalpy
13. H2O (g) + ; Δ H = 502 KJ/mol
+ ; Δ H = 427 KJ/mol
Average Bond Energy =
502 + 427
2
= 464.5 KJ/mol
Bond Enthalpy
14. ⮚ In the Lewis description of covalent bond, the Bond Order is given by
the number of bonds between the two atoms in a molecule.
Example:
⮚ Isoelectronic molecules and ions have identical bond orders.
Example: F2 and O2
-2
1) H2 bond order = 1
2) O2 bond order = 2
3) N2 bond order = 3
Bond Order
15. ⮚ Bond order ∝ bond enthalpy.
⮚ Bond order ∝
Bond Order
16. Sol:
CONCEPTUAL QUESTION
⮚ Arrange the following in order of decreasing bond angles CH4,
NH3,H2O, BF3, C2H2
C2H2(1800) > BF3(1200) > CH4(1090281) > NH3(1070) > H2O(104.50).
Bond Order
17. Which of the following has largest bond angle ?
a) NH3
b) PH3
c) AsH3
d) SbH3
18. The bond energy in halogens decreases in the order…
a) F2>Cl2>Br2>l2
b) Cl2>Br2>F2>l2
c) Cl2>F2>Br2>l2
d) F2>Br2>Cl2>l2
19. Carbon - Carbon bond length is least in…
1) C2H6
2) C6H6
3) C2H4
4) C2H2
22. 1. Attractive Force between:
(i) Electron
eA and nucleus HB.
(ii) Electron
eB and nucleus HA.
2. Repulsive Force
between:
(i) Electron
eA and nucleus eB.
(ii) Electron
HB and nucleus HB.
Valence Bond Theory
28. Heitler & London concept
1. Atoms must come
closer for
overlapping
Valence Bond Theory
29. Heitler & London
concept
Extent of Overlapping ∝ strength of
bond
2. Orbital Configuration
A B
A B
b.
3.
a. 1
1
1
1 +
Valence Bond Theory
30. 1. Extent of overlapping depends upon
a. Nature of Overlapping orbitals
b. Nature of Overlapping
A. More closer the valence shells are to the nucleus,
more will be the overlap.
B. Directional orbitals show more overlap.
Pauling and Slater’s Extension
Valence Bond Theory
31. Sigma Bond:
⮚ This type of covalent bond is formed by the end-on-end overlap of
bonding orbitals along the internuclear axis.
⮚ This is called as head on overlap or axial overlap.
Types Of Overlapping
32. X +↑
Y
+ +
↑
Z
Y
Z
X
+ + X
Z
Y
s – s overlap
s orbital s orbital
S-S overlapping: In this case, there is overlap of two half filled S-orbitals
along the internuclear axis.
S-S overlapping
Types Of Overlapping
33. S-P overlapping: This type of overlap occurs half filled S-orbital of one
atom & half filled P-orbitals of another atom.
p orbital
+↑ +
s orbital
s – p orbital
+
(Sigma bonds)
S-P overlapping
Types Of Overlapping
34. P-P overlapping: This type of overlapping takes place between half filled
P-orbitals of the two approaching atoms along their direction.
X
Y
Z
p subshell
+
↑
p orbital
↑
p orbital p – p overlap
(sigma bond
Axial overlap)
p orbital
↑
p – p overlap
(pi bond
Lateral
overlap)
p orbital
↑
+ P-P
overlapping
Types Of Overlapping
40. Sigma (σ) bond Pi (π) bond
1. It is formed by end to end overlapping
of half filled atomic orbitals.
1. It is formed by sidewise overlapping of
half filled p-orbitals only.
2. Overlapping takes place along
internuclear axis.
2. Overlapping takes place perpendicular
to internuclear axis.
3. The extent of overlapping is large and
bond formed is stronger.
3. The extent of overlapping is small and
bond formed is weaker.
4. The molecular orbital formed as a
result of overlapping is symmetrical
about the internuclear axis.
4. The molecular orbital formed as a
result of overlapping consists of two
lobes above and below the internuclear
axis.
5. There is free rotation about σ bond
and no geometrical isomers are
possible.
5. There is no free rotation about π bond
and geometrical isomers are possible.
6. The bond can be present alone. 6. The bond is always formed in addition
to sigma (σ) bond.
7. s and p orbitals can participate in the
formation of σ bond.
7. Only p-orbitals participate in the
formation of π bond.
𝞼 Vs 𝛑 Bonds
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