Chemical Bonding

1
Chapter 1
Chemical Bonding
Sections below not examined, but must be
understood!! Section 1.1 Atoms and Electrons
1.2 Ionic Bonds 1.3 Covalent Bonds 1.4 Double
Bonds and Triple Bonds 1.8 Isomers and
Isomerism 1.10 Shapes of Some Simple Molecules
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1.5 Polar Covalent Bonds 1.11 Molecular Polarity
H - H no net polarization
homopolar or non-polar covalent bond
H - F polarization of bonding electrons
partial negative charge (?-) on F partial
positive charge (?) on H
? charge separation H - F bond - polar
covalent bond. HF - polar molecule

H - F
?H - F?-
Polarity dipole moment ? - measure of
charge separation.
? relate to charge and bond length
? e.d Debye units e charge electrostatic
units (esu) d distance in cm. 1 Debye
Unit 1.0 x 10-18 esu.cm
3
1.11 Molecular Polarity Dipole Moment
Examples

• Dipole
• HF 1.9 Debye (1.9 x 10-18 esu.cm)
• HCl 1.1
• HBr 0.9
• HI 0.4

• Dipole Bond length
• CH3F 1.82 D C-F 1.39 x 10-8 cm
• CH3Cl 1.94 C-Cl 1.79
• CH3Br 1.79 C-Br 1.93
• CH3I 1.64 C-I 2.14

• Charge separation very important in organic
  chemistry ? we shall see how it determines
  reactivity later.

• In organic chemistry, use concept of
  electronegativity to provide qualitative guide
  to charge separation

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1.11 Molecular Polarity Electronegativity

• electronegativity

the ability of an atom in a bond to
attract electrons to itself Pauling
Electronegativity scale (Table 1.2 p. 8)
H 2.1 F 4.0 Cl 3.0 Br 2.8 I 2.5
Direction of polarization?

F Cl ?

I Cl ?
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1.11 Molecular Polarity Electronegativity (cont.)
H 2.1 C 2.5 F 4.0 Cl 3.0 Br 2.8 I 2.5
Electronegative element atom of element in a
compound with higher electronegativity than
carbon.
Metals donate electrons electropositive
lower electronegativity than carbon.
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1.11 Molecular Polarity Dipole Moment

• Compounds containing carbon bonded to sulfur or
  iodine may
• still possess dipole moments

• Table provide qualitative guide only to relative
  differences
• between related series of compounds.

• It is not possible to estimate magnitude of
  dipole moments
• from table

Just need to recognize direction of dipole!!
1.47 D
0.58 D
0.61 D
1.19 D
1.85 D
0.97 D
1.30 D
1.50 D
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1.11 Molecular Polarity Dipole Moment (cont.)
Presence of polar bond does not necessarily
confer a dipole moment
molecule may have net dipole moment of zero
(dipole vector quantity)
H2O
NH3
CO2
CH4
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1.11 Molecular Polarity Dipole Moment (cont.)
Which of the following molecules would you expect
to have a dipole moment? What is the direction
of the dipole?
C2H4
HCN
CO
CH2O
CH2Cl2

• Draw out structural formula, inspect atom
  electronegativities
• (Table 1.2, p. 8), and then draw bond dipoles
• Sum individual dipoles to provide net dipole

? 1.10 Shapes of Simple Molecules
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1.11 Molecular Polarity Electrostatic Potential
Map
- display 3D representation of 'charge'
distribution in an organic molecule very useful
for appreciating chemical reactivity!
Red negative Blue positive
HCl
Ethylene
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1.10 Shapes of Simple Molecules 1.12-1.15
Orbital Hybridization

• Express bond lengths in ?ngstrom (1 Å 10-8 cm)
  or
• picometres (1 pm 10-10 cm)
• Express bond angle in degrees ()
• Use thickened wedges, dashed wedges, and even
  lines to
• indicate 3D structure

Methane
1.34 Å
Ethylene
109.5
117.2
1.10 Å
1.10 Å
121.4
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1.12-1.15 Orbital Hybridization Valence Shell
Electron Pair Repulsion (VSEPR)
The most stable arrangement of groups attached to
a central atom is the one which has the maximum
separation of electron pairs bonded or
non-bonded
Ammonia
Water
Methane
Trigonal pyramidal'
Bent
tetrahedral arrangement of electron pairs
H-N-H 107?
H-O-H 105?
H-C-H 109.5?
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1.12-1.15 Orbital Hybridization VSEPR (cont.)
The most stable arrangement of groups attached to
a central atom is the one which has the maximum
separation of electron pairs bonded or
non-bonded
Cl-Be-Cl 180 ?
F-B-F 120 ?
linear
trigonal planar
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1.12-1.15 Orbital Hybridization i.
sp3-hybridization
Methane is tetrahedral!
However. configuration of orbitals on carbon ?
109.5
1s
2s
2p
1.10 Å
x
y
z



-
-
-
yz nodal plane
xz nodal plane
xy nodal plane
-
-
-
px-orbital boundary surface
py-orbital boundary surface
pz-orbital boundary surface
The px, py, and pz orbitals are capable of
bonding, but are orthogonal!
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1.12-1.15 Orbital Hybridization
sp3-hybridization (cont.)
Methane is tetrahedral!
Orbital hybridization
Equal in energy
Four singly occupied sp3 hybrid orbitals
For single carbon atom, promote electron from
2s to vacant 2p orbital
Hybridize the three 2p and one 2s orbitals
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1.12-1.15 Orbital Hybridization
sp3-hybridization (cont.)
2p
109.5
1.10 Å
2sp3
2s
The four sp3 hybrid orbitals towards the vertices
of tetrahedron
Other elements in the same period
H
H
H
H
H
H
H
H
H
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1.12-1.15 Orbital Hybridization ii.
sp2-hybridization
For single carbon atom, promote electron from
2s to vacant 2p orbital
Hybridize two 2p orbitals and one 2s
orbital
1 ? bond
3 ? bonds
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1.12-1.15 Orbital Hybridization iii.
sp-hybridization
For single carbon atom, promote electron from
2s to vacant 2p orbital
Hybridize one 2p and one 2s orbital
2 ? bond
2 ? bonds
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1.9 Resonance
III
A mule cross breed between a horse and a donkey
- a 'hybrid'
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1.9 Resonance

• Two or more structures of single molecule or ion
  with different
• bond types, overall electron counts generally
  obey the
• octet rule - resonance contributors.

ii. Relate by resonance arrow
iii. Structures differ through distribution of
electron pairs, bond types each structure
same number of unpaired, bonding electrons, same
net charge.
Nitrous acid
Nitrate ion
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1.9 Resonance (cont.)

• Two or more structures of single molecule or ion
  with different
• bond types, overall electron counts may obey
  the
• octet rule, but many cases not - resonance
  contributors.

ii. Relate by 'resonance arrow'
iii. Structures differ through distribution of
electron pairs, bond types each structure
same number of unpaired, bonding electrons, same
net charge.
Which of the following are resonance contributors?
X
?
24e
26e
X
unpaired electrons
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1.9 Resonance (cont.)
The structure of the molecule is not adequately
represented by an individual resonance
contributor, but is a hybrid of the resonance
contributors.
Consider acetic acid and acetate ion
1.25 ?
1.21 ?
1.25 ?
1.36 ?
- For acetate, resonance contributors differ in
position of charge, and double and single
bonds, but are of equal energy.
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1.9 Resonance (cont.)
Note that individual resonance contributors to
the structure may not be equal in in energy if
- one contributor has charge separation
A
B
B higher energy through charge separation
minor contributor
- resonance contributors of a charged ion have
the formal charge on different atoms
A lower energy - negative charge on more
electronegative atom major contributor
A
B
It is common practice to draw one resonance
contributor, with the understanding that other
contributors are implied.
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1.9 Resonance (cont.)
Benzene Michael Faraday 1825 Mitscherlich 1834
-C6H6 Kekulé, Ladenburg
Resonance Contributors equal in energy
C-C 1.40 Å C-H 1.08 Å ?CCC 120º
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Revision from A-Level Atomic Orbitals and
Quantum Numbers
Please Read and Understand!!
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Atomic Orbitals and Quantum Numbers
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Atomic Orbitals and Quantum Numbers (cont.)
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Size, Shape and Spatial Orientation of Atomic
Orbitals
Boundary surface encloses 90-95 of total
electron density. Orbital is approximated by
volume enclosed by boundary surface.

Probability distribution density - of
electron in 1s orbital wave function defining
distribution is assigned as positive

-


-

1s
2s
3s
Probability
Distance
Nucleus
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Size, Shape and Spatial Orientation of Atomic
Orbitals (cont.)
p-orbitals

-
yz is nodal plane
-
-

px-orbital boundary surface

xz is nodal plane
-
-
xy is nodal plane
py-orbital boundary surface
-
pz-orbital boundary surface
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Electron Occupancy of Atomic Orbitals
Need to consider a fourth quantum number
describing spin of electron
m may have value of either ½ or ½.
Total of four quantum numbers n, l, ml, m
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Electron Occupancy of Atomic Orbitals (cont.)
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Electron Occupancy of Atomic Orbitals (cont.)
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Electron Occupancy of Atomic Orbitals (cont.)