Chemical Bonding 3 POLAR BONDS
University of Lincoln
presentation
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Definitions…
• A HOMONUCLEAR BOND is a bond
between two identical atoms
• A HETERONUCLEAR BOND is a bond
between different atoms
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Homonuclear & Heteronuclear
bonds
Ethane (C2H6)
Homonuclear bonds
Hydrazine (N2H4)
Hydrogen peroxide
(H2O2)
Hetronuclear bonds
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Determining Bond Energies
• Consider the 2 homonuclear diatomics H2 and
F2
• The bond energy of H–F would be expected to
be the mean of the bond energies of H–H and
F–F
• Is this right?
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Bond Energies
Bond Dissociation Energy (kJmol-1)
X
Y
X–X
Y–Y
½ (X–X +
Y–Y)
H
F
436
159
298
Exptl
X–Y
570*
H
Cl
436
242
339
432*
H Br
436
193
315
366*
H
436
151
294
298
I
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Anomalous Bond Energies
Molecule
Expected Measured Bond
E
Bond
Energy (kJmol
H–F
Energy
(kJmol-1)
298
1)
570
272
H–Cl
339
432
93
H–Br
315
366
51
H–I
294
298
4
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?
Why are some heteronuclear
bonds much stronger than
expected?
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SYMMETRICAL BONDS
In a HOMONUCLEAR diatomic molecule, the
electrons within the bond are shared equally
between the two atoms – a symmetrical bond:
The electrons sit in molecular orbitals which lie
EQUI-DISTANT from each atom
Energy
σ*(2s)
2s
2s
Li
Li
σ*(2s)
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ASYMMETRICAL BONDS
• In a HETERONUCLEAR diatomic molecule, the
electrons within the bond are NOT always shared
equally between the two atoms – an
asymmetrical bond.
• In an assymetrical bond, the electrons sit closer to
one atom than the other, leading to a POLAR
BOND:
+
–

H–F

The electrons are sitting closer to the F atom
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Why does this happen?
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Electronegativity
Pauling defined ELECTRONEGATIVITY
as:
“the power of an atom in a molecule to
attract electrons to itself”
This is an atomic property, but only
applies when the atoms are in a
bond
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Electronegativity
H
2.2
Li
Be
B
C
N
O
F
1.0
1.6
2.0
2.6
3.0
3.4
4.0
Na
Mg
Al(III)
Si
P
S
Cl
0.9
1.3
1.6
1.9
2.2
2.6
3.2
K
Ca
Ga(III)
Ge(IV)
As(III)
Se
Br
0.8
1.0
1.8
2.0
2.2
2.6
3.0
Rb
Sr
In(III)
Sn(IV)
Sb
Te
I
0.8
0.9
1.8
2.0
2.1
2.1
2.7
Cs
Ba
Tl(III)
Pb(IV)
Bi
Po
At
0.8
0.9
2.0
2.3
2.0
2.0
2.2
The higher the electronegativity, the stronger the
‘pulling’ power of the atom within a bond
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So, why are some heteronuclear bonds much
stronger than expected?
…When electrons are held tightly by
an atom in a bond, due to the high
electronegativity of that atom, the
bond is much harder to break
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Examples of Polar Bonds
-
+
F H
H
+
H O –
-
Cl
+
The slight charges on each end of
the molecule lead to electrostatic
attraction between adjacent
molecules – HYDROGEN
BONDING
+
H
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Definition…
• A HYDROGEN BOND is an interaction between
a hydrogen atom attached to an electronegative
atom, and an electronegative atom which
possesses a lone pair of electrons
The strongest hydrogen bonds involve the first
row elements F, O or N
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HYDROGEN BONDING (
H–F
H–F
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)
25
20
15
PH3
AsH3
Molecule
SbH3
400
Boiling point (K)
Boiling point (K)
NH3
Boiling point (K)
ΔvapH/kJ mol-1
Hydrogen bonding affects the physical
properties of molecules with polar bonds
350
300
250
200
150
H2O
H2S
H2Se
Molecule
H2Te
270
250
230
210
190
170
150
NH3
PH3
AsH3
SbH3
HF
HCl
HBr
HI
Molecule
350
300
250
200
150
Molecule
NH3, H2O and HF all have anomalously HIGH boiling points,
since extra energy is needed to break the hydrogen bonds
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?
Can Molecular Orbital Theory
account for polar bonds?
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A quick recap…
ATOMIC
Orbitals
H + H
MOLECULAR
Orbitals
H2
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F2
Electronic configuration of 9F is:
1s2 2s2 2p5 (9 electrons)
The F atom needs 1 more electron to give it a full valence shell (8 outer
electrons)– it does this by forming a single covalent bond (in this
case with another F atom)
Hence, we know we have a single bond in F2: F–F
F F
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BUT we know that the F–F molecule has 18
electrons (2 x 9)
How can we arrange 18 electrons in molecular
orbitals and end up with only ONE bond?
SOLUTION:
•For every bonding orbital there must be an ‘antibonding orbital’
•An electron in a bonding orbital is cancelled out
by an electron in an anti-bonding orbital
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Consider the MO
diagram of F2
Energy
σ*(2pZ)
2p
π*(2px)
π*(2py)
π(2px)
π(2py)
2p
σ (2pZ)
σ*(2s)
2s
2s
F
σ (2s)
F
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Heteronuclear Diatomic molecule MO
Energy
σ*(2s)
2s
2s
X
σ*(2s)
Y
Homonuclear MO diagrams are symmetrical. Heteronuclear MOs are asymmetrical – the
energies of equivalent atomic orbitals are DIFFERENT
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LiH molecule
Energy
σ*(2s)
2s
2s
Li
σ*(2s)
H
Only valence orbitals shown. The 1s (H) and 2s (Li) overlap to form the  and
* molecular orbitals
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HF
Energy
σ*
1s
Non-bonding
2p
σ
Non-bonding
2s
H
HF
F
The 2pz(F) can overlap with the 1s(H). The orbitals that do not overlap
form NON-BONDING MOs
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x
Bonding
Z
1s
2pz
H
F
Z
AntiBonding
1s
2px
H
F
The 1s orbital on the H overlaps with the 2pz on the F to form a -bond. No
overlap can occur between the 1s and the 2px or 2py, as these are pointing in
the wrong direction
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HF
+
H–F
Energy
σ*
1s
Non-bonding
-
2p
σ
Non-bonding
2s
H
HF
F
The electrons are sat closer to the F atomic orbitals than the H atomic orbitals.
Therefore it is predicted that the H–F bond would be POLAR
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LiF
Energy
σ*
+
-
Li–F
2s
Non-bonding
2p
σ
Non-bonding
2s
Li
LiF
F
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Hence, the MO theory can
predict POLAR bonds
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Summary
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What you should know…
• Difference between homonuclear and
heteronuclear bonds
• Explain why some heteronuclear bonds are harder
than expected to break
• How the presence of hydrogen bonding in
molecules affects some of their physical
properties, like boiling points
• How to draw the MO diagram of a heteronuclear
diatomic molecule, and understand how bonding,
anti-bonding and non-bonding orbitals are
formed
• Use the MO diagram to determine whether the
bonding is likely to be polar
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Definitions…
• Homonuclear bond
• Heteronuclear bond
• Polar bond
• Hydrogen bond
• Electronegativity
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Acknowledgements
•
•
•
•
•
•
•
JISC
HEA
Centre for Educational Research and Development
School of natural and applied sciences
School of Journalism
SirenFM
http://tango.freedesktop.org
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