Class 11 chemical bonding notes ppt part I.pdf

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1. Chemical bonding and
Molecular structure
Class XI
Prepared by
Dr. Tarang Tomar
PGT chemistry

2. What is a chemical bond ?
The attractive forces which holds various constituents such as atoms
molecules and ions together in different chemical species is called a
Chemical bond.
Basically it is the same kind of electrostatic attraction that binds the
electron of an atom to its positively charged nucleus to form a
molecule. This process is accompanied by decrease in energy.
Decrease in energy ∝ strength of the bond
Therefore, molecules are more stable than atoms.

8. Lattice energy:
Amount of energy released during the formation of
1 mole of ionic crystal from its constituent ions.
Na
+
+ Cl
-
NaCl L.E. = -ve
Amount of energy required to dissociate 1 mole of ionic crystal into its
constituent ions.
NaCl Na
+
+ Cl
-
L.E. = +ve

14. Questions based on Lattice
energy
Trick
L.E ∝charge
L.E. ∝1/size
Q1. Which one is having more lattice energy?
a.NaF, MgF2, AlF3
b.Na2O, MgO, Al2O3
c.Li2O, Li3N
d.NaCl, KCl
e.NaF, NaCl, NaBr
f.NaF, MgCl2

19. Cu
+
, Ag
+
, Au
+
,
Zn
2+
, Cd
2+
, Hg
2+
Follow PNGC

24. Covalent bonds have several characteristics, including:
•Low melting and boiling points:Covalent compounds have low melting
and boiling points because of the relatively weak forces between their
particles.
•
•Poor conductors:Covalent compounds are poor conductors of electricity
and heat in all states (solid, molten, or aqueous).
•
•Insoluble in water:Covalent compounds are generally insoluble in water,
but they can dissolve in organic solvents.
•
•Soft solids, liquids, or gases:Covalent compounds can exist as soft
solids, liquids, or gases.
•
•Brittle:Covalent compounds are brittle solids
Characteristics of Covalent compounds

27. Lewis Dot Structure

28. Kossel Lewis approach to
chemical bonding
Lewis postulated that atoms achieve a stable octet when linked via
chemical bonds.
In case of bonds formed from H2 , F2 etc. the bond formed from
sharing of electrons between the atoms. In this case each atom
attains a stable outer octet of electrons.
Lewis symbols: In the formation of a molecule only outer electrons or
group valence of the electrons take part in chemical bonding and
hence are called valence electrons.
In electrons of the s, p, and f orbital's do not take part in chemical
bonding. Only ‘d’ orbital's take part.

29. Lewis symbols
Lewis symbols as shown are given below:
Lewis symbols only show the group valence or electrons that take part in
chemical bonding and hence it is called valence electrons.
Significance of Lewis Symbols : The
number of dots around the symbol represents
the number of valence electrons. This number
common or group valence of the element. The
group valence of the elements is generally
either equal to the number of dots in Lewis
symbols or 8 minus the number of dots or
valence electrons.

31. Trick for Calculations:-
Total electron (Q) = V.E of all atoms + (-vecharge) –(+vecharge)
Bond Pair of electron (B.P e
-
) = 2 x no. of bonds
Lone Pair of electron = Q -B.P e
-
(i.e. (L.P e
-
) or = bond or ≡ bond)

32. Draw the lewisdot structure of the following
compounds:-
1. H
2
2. O
2
3. H
3O
+
4. NH
4
+
5. NO
3
-
6. O
3
7. CO
3
2-
8. NO
2
-
9. SO
4
2-
10. SO
2
11. SO
3
12. PO
4
3-
13. CO

33. # Formal Charge (F.C.):-
Formal charge is the charge assigned to an atom in a
molecule
Formula / Trick :-
F.C. = V –L –
??????
2
Where, V = Valence electron on atom
L = Lone pair on atom
B = Bonding electron on atom

35. Calculate the formal charge of each
atom in a molecule
1.NO
3
-1
2.SO
3
-2
3.NO
2
-1
4.COCl
2
5.CO
6.BH
3
7.N
2O
4

36. Limitations of the octet rule.
The incomplete octet of the central atom
In some compounds, the number of
electrons surrounding the central atom is less
than eight. This is especially in the case of
elements having less than 4 valence
electrons. Eg: LiCl , BeH2.

37. The odd electron molecules:
In molecules with an odd number of electrons like nitric oxide, NO
and nitrogen dioxide, NO2, the octet rule is not satisfied for all the
atoms.
•The expanded octet:
Elements in and beyond the third period of the periodic table
have, apart from 3s and 3p orbitals, 3d orbitals also available for
bonding. In a number of compounds of these elements there are
more than eight valence electrons around the central atom. This is
termed as the expanded octet.
For example:-PCl5, SF6, H2SO4, SO3 etc

38. Coordinate Bond :-
•It is a covalent bond in which the shared electron pair from one
atom is known as coordinate bond.
•Necessary conditions for the formation of coordinate bond.
a. Octet of donor atom should be complete and should have at
least one lone pair of electron.
b. Acceptor atom should have a deficiency of at least 1 pair of
electron.
For eg.
c. Atom which provide electron pair for sharing is called donor and
other atom which accepts electron pair is known as acceptor.
•This is why the bond is known as dative bond.

43. Sigma and pi bond :-
Sigma and pi bondare two types of covalent bonds. It is a bond that is
formed by the mutual sharing of electrons so as to complete their octet
or duplet in the case of hydrogen, lithium and beryllium.
1. Sigma (??????) Bond
This type of covalent bond is
formed by the end-to-end (head-
on) overlap of bonding orbitals
along the internuclearaxis. This is
called head-on overlap or axial
overlap. This can be formed by
any one of the following types of
combinations of atomic orbitals.

44. 2. Pi () Bond
During the formation of Pi bonds, atomic orbitals overlap in such a way that
their axes remain parallel to each other and perpendicular to the internuclear
axis.

46. •All single bonds are ??????-bonds.
•Multiple bonds contain one??????-bond, and the rest are π-bonds.
•π-bond is never formed alone. First, a ??????-bond is formed, and then the
formation of the π-bond takes place. (Exception: C molecule contains
both π-bonds)
•A sigma bond is always stronger than api bond because the extent of
overlapping of atomic orbitals along the internuclearaxis is greater than
sideways overlapping.
•The electron cloud of ??????-bond is symmetrical about the internuclearaxis,
while that of π-bond is not.
•Free relation about a ??????-bond is possible, but that about a π-bond is not
possible.
•Thelesspi bonds,themorestablethecompoundis.
•Themorethenumberof pi bonds,thecompoundis morereactive.
Note:

47. Overlap of Atomic Orbitals
Theatomicorbitalsof2atomsoverlaponcetheyarenearertooneanother.
Theoverlappingofatomicorbitalscanhavepositive,negative,orzero
overlaps,relyinguponthesecharacteristics.Thefigurebeneathshowsthe
differentconfigurationsofthesandp-orbitalsthatleadtopositive,negative,
andzerooverlaps.
•Positiveatomicorbitaloverlap:Wheneverthetwoinvolvedatomicorbitals
phaseisidentical,positiveoverlaptakesplace.Bondsarecreatedasa
consequenceofthisoverlap.
•Negativeatomicorbitaloverlap:Negativeoverlapoccurswheneverthe
phasesoftheinvolvedatomicorbitalsopposeoneanother.Bondformation
doesn'ttakeplaceinthisinstance.
•Zerooverlapsofatomicorbital:Zeroatomicorbitaloverlapsoccurwhile
twointriguingorbitalsdonotoverlapwitheachotherinanorbital.Theorbital
overlapdiagramisshowninnextpage:-

48. Zero overlap
Zero overlap

49. Hybridization
Pauling introduced the concept of hybridisation.
According to him the atomic orbitals combine to form
new set of equivalent orbitals known as hybrid orbitals.
Unlike pure orbitals, the hybrid orbitals are used in bond
formation. The phenomenon is known as hybridisation
which can be defined as the process of intermixing of
the orbitals of slightly different energies so as to
redistribute their energies, resulting in the formation of
new set of orbitals of equivalent energies and shape. For
example when one 2s and three 2p-orbitals of carbon
hybridise, there is the formation of four new Sp
3
hybrid
orbital

64. Limitations of Valence Bond Theory
it fails to explain the shape and geometry of molecules.
It fails to explain magnetic behavior of covalent molecules.

65. Valence Shell Electron Pair
Repulsion Theory
•VSEPR Theory was suggestedby Sidgwick and
Powel[1940]
•It was developed by Gilllespeand Nyholmin 1957.
•Based on that in a polyatomic molecule the
direction bonds around the central atom depends
on the total number of Bonding &Non-bonding
electron pairs in its valance shell.

66. •The shape of the molecule is determined by
repulsions between all of the electron present in the
valance shell.
•Electron pairs in the valence shell of the central
atom repel each other and align themselves to
minimize this repulsion.
•Lone pair electrons takes up more space round the
central atom than a bondpair.
•Lone pair attracted to one nucleus, but bond pair is
shared by two nuclei.
•The minimum repulsions to the state minimum
energy and maximum stability of the molecule.
Postulates of VSEPR Theory

67. Repulsion strengths
Lone pair-Lone pair Lone pair-
Bond pair Bond pair-Bond pair
Nyholmnand Gillespie stated that the electron pairs
existing as l.p.causes greater repulsive interactions
as compared to bonded pairs

71. Limitations of VSEPR theory
It fails to predict the shapes ofisoelectronic
species(CH
4&NH
4
+
) and transition metal
compounds.
This model does not take relative sizes of
substituents.
Unable to explain atomic orbitalsoverlap.

72. Molecular Orbital Theory was proposed by Hund & Mulliken.
With the help of MOT, we can explain and understand those
things that VBT was unable to explain. (Exp. Paramagnetic nature
of O
2molecule. As per VBT, it should be Diamagnetic.)
The atomic orbital loses its identity during molecule formation (by
overlapping) and forms new orbitals called molecular orbitals.
Molecularorbitalformedbyoverlappingofatomicorbitalofsame
energy.
Electronsinamoleculeoccupymolecularorbitalsinaccordance
withAufbauprinciple,Pauli’sexclusionprincipleandHund’srule.
No.ofmolecularorbitalsformed=No.ofatomicorbitalsinvolved
inoverlapping.
Molecular Orbital Theory
Salient features of Molecular orbital theory:-

73. According to this theory, all the atomic orbitals of the
participating atoms gets disturbed when the concerned nuclei
approach each other.
Definition of Atomic orbitals:-
“ The region in space around the nucleus of an atom when
the probability of finding the electron density is maximum.”
Definition of Molecular orbitals:-
“ The region in space comprising the nuclei of the combining
atom around which there is maximum probability of finding
the electron density.”

74. Linearcombination of atomic
orbitals (L.C.A.O)
Molecular orbitals are formed by the linear
combination of the wave function of participating
atomic orbitals
They may combine either by addition or subtraction.
Let ψA and ψB represents the wave functions of the
two combining atomic orbitals A and B.
Ψ= ψA + ψB Bonding Molecular orbitals
Ψ
*
= ψA -ψB Antibonding Molecular orbitals

75. Ψ= ψA + ψB
a. Combination by addition
b. Combination by subtraction
Ψ
*
= ψA -ψB
Constructive interference
Destructive interference

77. Shapes of molecular orbitals
1. Combination between s-atomic orbitals
Node
Note :-Node means nodal plane where the probability of finding electron
density is nill.

78. 2. Combination between s and p-atomic orbitals

79. 3. Combination between p
x-atomic orbitals

80. 4. Combination between p
yor p
z-atomic orbitals

83. Energy Level Diagram of Homoatomicmolecules
Get repelled
by magnetic
fields
Get attracted
by magnetic
fields

85. M.O. energy level diagram for
homonuclear diatomic molecule.
1. Hydrogen molecule (H
2)
Electronic Configuration of H (1): 1s
1
Electronic Configuration of H
2(2): σ1s
2
Energy
A.O. of
H
A.O. of
H
M.O. of H
2
σ*1s
σ1s

86. Bond order in H
2Molecule = ½ (Number of electron in BMO -Number of
electron in ABMO )
= ½ (2-0)
= ½ x 2
= 1
Bond order in H
2= 1……………i.e .(H-H)
Thus,thebondorderinH
2moleculeis1.Itsuggestthatthereisasingle
bondpresentbetweenthetwoH-atomsinH
2molecule.Itisaofσtype.
H
2moleculeisdiamagneticbecausealltheelectronsarepaired.

87. 1a. Hydrogen ion (H
2
+
ion)
Electronic Configuration of H
2(2): σ1s
2
Electronic Configuration of H
2
+
(1): σ1s
1
Energy
A.O. of
H
A.O. of
H
M.O. of H
2
σ*1s
σ1s
Electronic Configuration of H (1): 1s
1
Bond order in H
2
+
ion = ½ (Number of electron in BMO -Number of electron in
ABMO )
= ½ (1-0) = ½

88. 2. Helium molecule (He
2)
Electronic Configuration of He (2): 1s
2
Electronic Configuration of He
2(4): σ1s
2
, σ*1s
2
Energy
A.O. of
He
A.O. of
He
M.O. of
He
2
σ*1s
σ1s

89. Bond order in He
2Molecule = ½ (Number of electron in BMO -Number of
electron in ABMO )
= ½ (2-2)
= ½ x 0
= 0
Bond order in He
2= 0……………… i.e.(Molecule is
unstable)
Thus,thebondorderinHe
2moleculeis0.ItsuggestHe
2moleculeisnot
stable.HenceHe
2moleculedoesnotexist.Heliumbeinginertgas
elementexistinatomicformonly.

90. 3. Lithium molecule (Li
2)
Electronic Configuration of Li (3): 1s
2
, 2s
1
Electronic Configuration of Li
2(6): σ1s
2
, σ*1s
2
, σ2s
2
Energy
A.O. of Li A.O. of Li
M.O. of Li
2
σ
1s
σ*
1s
σ
2s
σ*
2s
1s 1s
2s 2s
Bond order in Li
2Molecule = ½ (Number of
electron in BMO -Number of electron in
ABMO )
= ½ (4-2)
= ½ x 2
= 1
Thus,thebondorderinLi
2moleculeis1.It
suggestthatthereisasinglebond
presentbetweenthetwoLi-atomsinLi
2
molecule.Itisaofσtype.TheLi
2
moleculeisdiamagneticduetothe
presenceofpairedelectrons.
Bond order in Li
2= 1………
i.e.(Molecule is stable)

91. 4. Beryllium molecule (Be
2)
Electronic Configuration of Be (4): 1s
2
, 2s
2
Electronic Configuration of Be
2(8): σ1s
2
, σ*1s
2
, σ2s
2
,σ*2s
2
Energy
A.O. of Be A.O. of
Be
M.O. of Be
2
σ
1s
σ*
1s
σ
2s
σ*
2s
1s 1s
2s2sBond order in Be
2Molecule = ½ (Number of
electron in BMO -Number of electron in
ABMO )
= ½ (4-4)
= ½ x 0
= 0
Bond order in Be
2= 0………………
i.e.(Molecule is unstable)
Thus,thebondorderinBe
2moleculeis0.It
suggestBe
2moleculeisnotstable.Hence
Be
2moleculedoesnotexist.Beryllium
beinginertgaselementexistinatomic
formonly.

92. 5. Boron molecule (B
2)
Electronic Configuration of B (5):
1s
2
,2s
2
,2px
1
Electronic Configuration of B
2(10):
σ1s
2
, σ*1s
2
, σ2s
2
,σ*2s
2
,σ2px
2
Energy
A.O. of B A.O. of B
M.O. of B
2
σ
2s
σ*
2s
2s 2s
2px2py2pz 2px2py2pz
σ*2px
π*2py π*2py
σ2px
π2py π2py

93. Bond order in B
2Molecule = ½ (Number of electron in BMO -Number of
electron in ABMO )
= ½ (4-2)
= ½ x 2
= 1
Bond order in B
2= 1……………… i.e.(B-
B)
Thus,thebondorderinB
2moleculeis1.Itsuggestthatthereisasingle
bondpresentbetweenthetwoB-atomsinB
2molecule.Itisaofσtype.
TheB
2moleculeisdiamagneticduetothepresenceofpairedelectrons.

94. 6. Carbon molecule (C
2)
Electronic Configuration of C (6):
1s
2
,2s
2
,2px
1
,2py
1
Electronic Configuration of C
2(12):
σ1s
2
, σ*1s
2
, σ2s
2
,σ*2s
2
,σ2px
2
,π2py
1
,
π2pz
1
Energy
A.O. of C A.O. of
C
M.O. of C
2
σ
2s
σ*
2s2s 2s
2px2py2pz 2px2py2pz
σ*2px
π*2py π*2py
σ2px
π2py π2py

95. Bond order in C
2Molecule = ½ (Number of electron in BMO -Number of
electron in ABMO )
= ½ (6-2)
= ½ x 4
= 2
Bond order in C
2= 2………………
i.e.(C=C)
Thus,thebondorderinC
2moleculeis2.Itsuggestthattherearetwo
bondspresentbetweenthetwoC-atomsinC
2molecule.Outoftwo
bondsoneisσbond&anotherisbond.TheC
2moleculeisparamagnetic
duetothepresenceoftwounpairedelectrons.
σ
π
π

96. 7. Nitrogen molecule (N
2)
Electronic Configuration
of N (7):
1s
2
,2s
2
,2px
1
,2py
1
,2pz
1
Electronic Configuration
of N
2(14): σ1s
2
, σ*1s
2
,
σ2s
2
,σ*2s
2
,σ2px
2
,
π2py
2
,π2pz
2
Energy
A.O. of N A.O. of
N
M.O. of
N
2
σ
2s
σ*
2s
2s 2s
2px2py2pz 2px2py2pz
σ*2px
π*2py π*2py
σ2px
π2py π2py

97. Bond order in N
2Molecule = ½ (Number of electron in BMO -Number of
electron in ABMO )
= ½ (8-2)
= ½ x 6
= 3
Thus,thebondorderinN
2moleculeis3.Itsuggestthattherearethree
bondspresentbetweenthetwoN-atomsinN
2molecule.Outofthree
bonds,oneisσbond&therearetwo bond.TheN
2moleculeis
diamagneticduetothepresenceofpairedelectrons.
Bond order in N
2= 3……………… i.e.(N= N)
σ
π
π
π

98. 8. Oxygen molecule (O
2)
Electronic Configuration of O (8):
1s
2
,2s
2
,2px
2
,2py
1
,2pz
1
Electronic Configuration of
O
2(16): σ1s
2
, σ*1s
2
, σ2s
2
,
σ*2s
2
,σ2px
2
,π2py
2
,π2pz
2
,
π*2py
1
,π*2pz
1
,
Energy
A.O. of O A.O. of
O
M.O. of O
2
σ2s
σ* 2s
2s 2s
2px2py2pz 2px2py2pz
σ*2px
π*2py π*2py
σ2px
π2py π2py

99. Bond order in O
2Molecule = ½ (Number of electron in BMO -Number of
electron in ABMO )
= ½ (8-4)
= ½ x 4
= 2
Bond order in O
2= 2 (stable)……………… i.e.(O O)
Thus,thebondorderinO
2moleculeis2.Itsuggestthattherearetwo
bondspresentbetweenthetwoO-atomsinO
2molecule.Outoftwo
bondsoneisσbond&oneisπbond.TheO
2moleculeisparamagnetic
duetothepresenceoftwounpairedelectrons.

100. 8a. Oxygen molecule (O
2
+
ion)
Electronic Configuration of O (8): 1s
2
,2s
2
,2px
2
,2py
1
,2pz
1
Electronic Configuration of O
2(16): σ1s
2
, σ*1s
2
, σ2s
2
,σ*2s
2
,σ2px
2
,π2py
2
,π2pz
2
, π*2py
1
,
π*2pz
1
Electronic Configuration of O
2
+
(15): σ1s
2
, σ*1s
2
, σ2s
2
,σ*2s
2
,σ2px
2
,π2py
2
,π2pz
2
, π*2py
1
Bond order in O
2
+
ion = ½ (Number of electron in BMO -Number of electron
in ABMO )
= ½ (8-3)
= ½ x 5
= 2.5
Bond order in O
2
+
= 2.5………… i.e.(O= O)
Thus,thebondorderinO
2
+
ionis2.5.InO
2
+
ionthereisoneσbond,one
bond&onethreeelectronbond.
…
π
½
σ
π

101. 8b. Oxygen molecule (O
2
-
ion Superoxide ion)
Electronic Configuration of O (8): 1s
2
,2s
2
,2px
2
,2py
1
,2pz
1
Electronic Configuration of O
2(16): σ1s
2
, σ*1s
2
, σ2s
2
,σ*2s
2
,σ2px
2
,π2py
2
,π2pz
2
, π*2py
1
,
π*2pz
1
Electronic Configuration of O
2
-
(17): σ1s
2
, σ*1s
2
,σ2s
2
,σ*2s
2
,σ2px
2
,π2py
2
,π2pz
2
, π*2py
2
,
π*2pz
1
Bond order in O
2
-
ion = ½ (Number of electron in BMO -Number of electron
in ABMO )
= ½ (8-5)
= ½ x 3
= 1.5
Bond order in O
2
-
= 1.5………… i.e.(O -O)
Thus,thebondorderinO
2
-
ionis1.5.InO
2
-
ionthereisoneσbond&one
threeelectronbond.
…
σ

102. 8c. Oxygen molecule (O
2
--
ion peroxide ion)
Electronic Configuration of O (8): 1s
2
,2s
2
,2px
2
,2py
1
,2pz
1
Electronic Configuration of O
2(16): σ1s
2
, σ*1s
2
, σ2s
2
,σ*2s
2
,σ2px
2
,π2py
2
,π2pz
2
, π*2py
1
,
π*2pz
1
Electronic Configuration of O
2
--
(18): σ1s
2
, σ*1s
2
,σ2s
2
,σ*2s
2
,σ2px
2
,π2py
2
,π2pz
2
, π*2py
2
,
π*2pz
2
Bond order in O
2
-
ion = ½ (Number of electron in BMO -Number of electron
in ABMO )
= ½ (8-6)
= ½ x 2
= 1
Bond order in O
2
--
= 1………… i.e.(O -O)
Thus,thebondorderinO
2
--
ionis1.InO
2
--
ionthereisoneσbond.
σ
Bond Order
O
2
+
>O
2>O
2
-
>O
2
--
2.52 1.51

103. 9. Fluorine molecule (F
2)
Electronic Configuration of F (9):
1s
2
,2s
2
,2px
2
,2py
2
,2pz
1
Electronic Configuration of
F
2(18): σ1s
2
, σ*1s
2
, σ2s
2
,σ*2s
2
,
σ2px
2
,π2py
2
,π2pz
2
, π*2py
2
,
π*2pz
2
,
Energy
A.O. of F A.O. of F
M.O. of F
2
σ
2s
σ*
2s2s
2s
2px2py2pz 2px2py2pz
σ*2px
π*2py π*2py
σ2px
π2py π2py

104. Bond order in F
2Molecule = ½ (Number of electron in BMO -Number of
electron in ABMO )
= ½ (8-6)
= ½ x 2
= 1
Bond order in F
2= 1……………… i.e.(F-F)
Thus,thebondorderinF
2moleculeis1.Itsuggestthatthereisasingle
bondpresentbetweenthetwoF-atomsinF
2molecule.Itisaofσtype.
TheF
2moleculeisdiamagneticduetothepresenceofpairedelectrons.

105. 10. Neon molecule (Ne
2)
Electronic Configuration of F(9):
1s
2
,2s
2
,2px
2
,2py
2
,2pz
2
Electronic Configuration of
F
2(18): σ1s
2
, σ*1s
2
, σ2s
2
,
σ*2s
2
,σ2px
2
,π2py
2
,π2pz
2
,
π*2py
2
,π*2pz
2
,σ*2px
2
Energy
A.O. of
Ne
A.O. of
Ne
M.O. of Ne
2
σ
2s
σ*
2s
2s 2s
2px2py2pz 2px2py2pz
σ*2px
π*2py π*2py
σ2px
π2py π2py

106. Bond order in Ne
2Molecule = ½ (Number of electron in BMO -Number of
electron in ABMO )
= ½ (8-8)
= ½ x 0
= 0
Bond order in Ne
2= 0……………… i.e.(Molecule is unstable)
Thus,thebondorderinNe
2moleculeis0.ItsuggestNe
2moleculeisnot
stable.HenceNe
2moleculedoesnotexist.Neonbeinginertgaselement
existinatomicformonly.

109. For example:-
1. Hydrogen Bond in Water (H
2O)
A highly electronegative oxygen atom is
connected to a hydrogen atom in the water
molecule. The shared pair of electrons are
attracted to the oxygen atoms more, and
this end of the molecule becomes
negative, while the hydrogen atoms
become positive.

110. 2. Hydrogen Bond in Hydrogen Fluoride(HF)
A stronger-than-average hydrogen bond is created by hydrofluoric acid and is
known as a symmetric hydrogen bond. Formic acid can also make this type of
bond.

111. 3. Hydrogen Bond in Ammonia (NH
3)
Between the hydrogen in one molecule and the nitrogen in another, hydrogen
bonds are formed. Since each nitrogen has a single electron pair, the bond that
develops in the case of ammonia is relatively weak. Methylamine also has this
form of hydrogen bonding with nitrogen.

112. 4. Hydrogen Bond in Alcohol and Carboxylic Acid
A type of chemical molecule with a -OH group is alcohol. In most cases,
hydrogen bonding is easily generated if any molecule containing the hydrogen
atom is immediately coupled to either oxygen or nitrogen.
Hydrogen Bond in Alcohol

113. Hydrogen Bond in Carboxylic Acid
Strength of Hydrogen Bond:The hydrogen bond is a relatively weak one.
Hydrogen bonds have a strength that is halfway between weak van der Waals
forces and strong covalent bonds. The attraction of the shared pair of electrons,
and hence the atom’s electronegativity, determines the hydrogen bond’s
dissociation energy.

114. 1.Volatility –The boiling point of compounds incorporating hydrogen bonding
between distinct molecules is greater, hence they are less volatile.
2.Solubility –Because of the hydrogen bonding that can occur between water
and the alcohol molecule, lower alcohols are soluble in water.
3.Lower density of ice than water –In the case of solid ice, hydrogen bonding
causes water molecules to form a cage-like structure. In fact, each water
molecule is tetrahedrallyconnected to four other water molecules. In the solid
state, the molecules are not as tightly packed as they are in the liquid state. This
case-like structure collapses as ice melts, bringing the molecules closer together.
As a result, the volume of water reduces while the density increases for the
same quantity of water. As a result, at 273 K, ice has a lower density than water.
Ice floats because of this.
4.Viscosity and Surface Tension –Hydrogen bonding is found in compounds
that have an associated molecule. As a result, their flow becomes more
complicated. They have high surface tension and higher viscosity.
Properties of Hydrogen Bonding

115. Types of Hydrogen Bonding
1. Intermolecular Hydrogen Bonding
Intermolecular hydrogen bonding occurs
when hydrogen bonds are formed between
molecules of the same or distinct
substances. Hydrogen bonding in water,
alcohol, and ammonia, for example.

116. 2. Intramolecular Hydrogen Bonding
Intramolecular hydrogen bonding refers to hydrogen bonding that occurs within a
single molecule. It occurs in compounds with two groups, one of which has a
hydrogen atom linked to an electronegative atom and the other of which has a
highly electronegative atom linked to a less electronegative atom of the other
group. The link is created between the more electronegative atoms of one group
and the hydrogen atoms of the other group.
For eg. o-nitrophenol, o-hydroxy benzoic acid etc.