HOMONUCLEAR COVALENT BONDS
University of Lincoln
presentation
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Chemical Bonds
A CHEMICAL BOND joins atoms together
There are 4 types of chemical bond:
•
•
•
•
COVALENT BONDS
Ionic bonds
Coordinate bonds
Metallic bonds
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Homonuclear Covalent Bonding
What you need to know…
• Covalent bond formation
• Bond length
• Bond energy
• Bond order
• Relationship between bond length, bond
energy and bond order
• Trends in the periodic table
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Definitions…
• A MOLECULE is a discrete neutral species
resulting from the formation of a covalent
bond or bonds between two or more atoms
• A HOMONUCLEAR BOND is a covalent
bond between 2 identical atoms
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Covalent Homonuclear
Molecules
Hydrogen (H2)
Oxygen (O2)
Ozone (O3)
Phosporous (P4)
Iodine (I2)
Sulphur (S6)
Examples of covalent homonuclear molecules
Sulphur (S8)
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Molecules with Homonuclear Bonds
Ethane (C2H6)
Hydrazine (N2H4)
Hydrogen peroxide
(H2O2)
Molecules with one homonuclear bond
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Making a Covalent Bond –
sharing valence electrons
In order to share valence electrons, 2 atoms have
to come into close contact with each other
H
H
1s1
1s1
2 hydrogen atoms
H H
He
1s2
1 hydrogen molecule, H2
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Bringing 2 atoms together is not easy –
there are FOUR forces in play…
+
–
ATOM A
+
ATOM B
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The Four Forces
Internuclear separation
(2)
+
-
(3)
(1)
+
(4)
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How close do the atoms have to be to
form a bond?
r
The VAN DER WAALS RADIUS (rv) of an atom X is measured as half
of the distance of closest approach of 2 NON-BONDED atoms of
X
The COVALENT RADIUS (rcov) of an atom X is taken as half of the
internuclear distance (r) in a HOMONUCLEAR X–X bond.
The internuclear distance (r) in a bonded pair of atoms is called the
BOND LENGTH
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Non-bonded vs Bonded Radii
Element
Van der Waals
radius (pm)
NON-BONDED
Covalent X–X
radius (pm)
BONDED
Covalent Bond
Length (pm)
(2 x rcov)
H
120
37
74
B
208
88
176
C
185
77
154
Si
210
118
236
N
154
75
150
O
140
73
146
S
185
103
206
F
135
71
142
Note: the internuclear distance is SMALLER when atoms
are bonded together
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…Hence, atoms must overlap to
form a bond
Bond length
Non-bonded atoms –
NO OVERLAP of
atomic orbitals
Bonded atoms –
OVERLAP of atomic
orbitals
The bigger the overlap, the SHORTER the bond.
The shorter the bond, the STRONGER it is.
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Bond Energy
Sometimes called the BOND ENTHALPY
The BOND ENERGY is the amount of energy
required to break a bond:
H–H
2H
The bond energy is, therefore, a
measure of how strong a bond is:
The larger the bond energy, the STRONGER the bond
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Breaking Bonds…
Breaking the C-C
bond produces two
radicals
C2H6
2CH3·
Breaking the S-S
bond opens up the
ring structure
S6
·S-S-S-S-S-S·
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Bond Order
Type of Bond
Name of Bond
Bond Order
X–X
Single
1
X=X
Double
2
X≡X
Triple
3
The larger the bond order, the STRONGER the bond
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Some Bond Energies
Bond
H–H
C–C
C=C
C≡C
N–N
N=N
N≡N
P–P
P≡P
Bond Energy
(kJmol-1)
436
346
598
Bond
O–O
O=O
S–S
Bond Energy
(kJmol-1)
146
498
266
813
159
400
S=S
F–F
Cl–Cl
425
159
242
945
200
Br–Br
I–I
193
151
490
Group 17
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Bond Energy & Bond Length
F–F
Bond
Energy
(kJmol-1)
159
Bond
Length
(pm)
141
300
250
200
150
Cl–Cl
242
199
100
Br–Br
193
228
50
I–I
151
267
0
F-F
Cl-Cl
Br-Br
I-I
The shorter the bond, the higher the bond energy
F is anomalous due to its small size. Bond energy would be
expected to be ~275 kJmol-1
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Adjacent Lone Pair Effect
F F
Because F is a small atom
(look at its position on the
Periodic Table – it is the
smallest of the 1st row
elements) its valence electrons
are very close and tend to
repel each other. The two
atoms are forced apart and
the bond is weakened
veryclose
This anomalous
behaviour is
common in 1st row
elements,
particularly, N, O
and F
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Group Trends in Homonuclear
Single Bond Energies
Bond Energy (kJmol-1)
350
300
250
Group14
Group 15
Group 16
Group 17
200
150
100
50
0
1st row
2nd row 3rd row
4th row
Note the anomalous behaviour of N–N, O–O and F–F. Group 14
show the expected trend
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Formation of Multiple bonds
N
Could make a double or a triple bond. A
triple bond would be stronger (sharing
all three unpaired electrons with
another atom)
O
Can only
make a
single bond
F
Could make a
double bond
(sharing both of
its unpaired
electrons with
another atom)
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Bond Energies for X2 Molecules in
Group 15 (in their natural state)
N2 has very high
bond
energy…why?
750
Bond
Ene rgy
(KJm ol-1)
500
Bond Energy (kJmol-1)
1000
250
0
N
P
As
Sb
Bi
Ele m e nt
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Formation of the N2 Molecule
N
N is small enough to overlap with
another N atom sufficiently to share all
three of its unpaired electrons and make
a very strong TRIPLE BOND
N N
LINEAR Molecule
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Other elements in Group 15…
• P, As, Sb and Bi are TOO BIG to form
multiple bonds – they can’t get close
enough to overlap sufficiently
• These elements form SINGLE BONDS with
three other atoms forming TETRAHEDRAL
molecules
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…Nitrogen forms a triple bond
Other elements in Group 15 can only form single bonds
P
2 X
X X
4 X
P
P
P
X=N
X = P, As, Sb or Bi
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280
1000
260
900
240
800
Bond enthalpy kJ mol-1
X - X bond distance in X2 molecule/pm
Periodic Trends in Bond Length,
Bond Energy & Bond Order
220
200
180
160
140
700
600
500
400
300
200
120
100
100
0
Li
B
C
N
X-X bond distances
O
F
Li
B
C
N
O
X-X bond dissociated
enthalpy for X2
molecules containing
the first row elements
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F
Bond Orders of the 1st Row
Elements
Homonuclear
Diatomic
B–B
Bond Order
C=C
2
N≡N
3
O=O
2
F–F
1
1
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Summary
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Things to Remember…
1. The covalent bond is formed by
overlapping atomic orbitals
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2. Bond Order (single, double, triple)
3. Bond Energy (energy to break bond, kJmol-1)
(measure of bond strength)
4. Bond Length (internuclear distance, pm)
5. Trends in the periodic table…
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TRENDS:
– Bond energy increases as bond order increases
– Bond length decreases as bond order increases
– Bond energy decreases as bond length
increases
The shorter the bond, the stronger it is
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Definitions
Molecule
Homonuclear bond
van der Waals radius
Covalent radius
Bond length
Bond energy
Bond order
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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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