36
ATOMIC STRUCTURE AND BONDING
the two atoms. Covalent bonds may, therefore,
possess a charge imbalance, with one of the atoms
taking more than its share of the electrons. This
is referred to as bond polarity. An atom that is
more electronegative than carbon will thus polarize
the bond, and we can consider the atoms as being
partially charged. This is indicated in a structure by
putting partial charges (δ+ and δ−) above the atoms.
It can also be represented by putting an arrowhead
on the bond, in the direction of electron imbalance.
Alternatively, we use a specific dipole arrow at the
side of the bond.
C Br
bromine is more
electronegative
than carbon
d+ d−
C
Br
C Br
partial
charges
dipole on
bond
dipole arrow
In general, electronegativities increase from left to
right across the periodic table, and decrease going
down a particular column of the periodic table.
The relative electronegativities of those atoms most
likely to be found in typical organic molecules
are included in Table 2.2. The numbers (Pauling
electronegativity values) are on an arbitrary scale
from Li = 1 to F = 4.
From the sequence shown, it is readily seen
that hydrogen and carbon are among the least
electronegative atoms we are likely to encounter in
organic molecules. The relatively small difference
in electronegativities between hydrogen and carbon
also means there is not going to be much polarity
associated with a C–H bond. Most atoms other than
hydrogen and carbon when bonded to carbon are
going to be electron rich; therefore, bonds may
Table 2.2 Pauling electronegativity values
H
2.1
Li
Be
B
C
N
O
F
1.0
1.6
2.0
2.5
3.0
3.5
4.0
Na
Mg
Al
Si
P
S
Cl
0.9
1.2
1.5
1.8
2.1
2.5
3.0
K
B r
0.8
2.8
I
2.5
display considerable polarity. This polarity helps us
to predict chemical behaviour, and it is crucial to our
prediction of chemical mechanisms.
C O
C N
d+ d−
d+ d−
polarity in
and
bonds
C O
C N
We must also modify our thinking of bonding
as being simply ionic (where there is transfer of
electrons between atoms) or covalent (where there
is equal sharing of electrons). These represent two
extremes, but bond polarity now provides a middle
ground where there is sharing of electrons, but an
unequal sharing.
Bond polarity in a molecule can often be measured
by a dipole moment, expressed in Debye units
(D). However, the physical measurement provides
only the overall dipole moment, i.e. the sum of the
individual dipoles. A molecule might possess bond
polarity without displaying an overall dipole if two
or more polar bonds are aligned so that they cancel
each other out. The C–Cl bond is polar, but although
chloroform (CHCl 3 ) has a dipole moment (1.02 D),
carbon tetrachloride (CCl 4 ) has no overall dipole.
Because of the tetrahedral orientation of the dipoles
in carbon tetrachloride, the vector sum is zero.
H
Cl
C
Cl
Cl
overall dipole
Cl
Cl
C
Cl
Cl
tetrahedral orientation of
dipoles means vector sum
is zero
Polarization in one bond can also influence the
polarity of an adjacent bond. Thus, in ethyl chloride,
the polarity of the C–Cl bond makes the carbon more
positive (δ+); consequently, electrons in the C–C
bond are drawn towards this partial positive charge.
The terminal carbon thus also experiences a partial
positive charge, somewhat smaller than δ+ and so
depicted as δδ+.
ATOMIC STRUCTURE AND BONDING
the two atoms. Covalent bonds may, therefore,
possess a charge imbalance, with one of the atoms
taking more than its share of the electrons. This
is referred to as bond polarity. An atom that is
more electronegative than carbon will thus polarize
the bond, and we can consider the atoms as being
partially charged. This is indicated in a structure by
putting partial charges (δ+ and δ−) above the atoms.
It can also be represented by putting an arrowhead
on the bond, in the direction of electron imbalance.
Alternatively, we use a specific dipole arrow at the
side of the bond.
C Br
bromine is more
electronegative
than carbon
d+ d−
C
Br
C Br
partial
charges
dipole on
bond
dipole arrow
In general, electronegativities increase from left to
right across the periodic table, and decrease going
down a particular column of the periodic table.
The relative electronegativities of those atoms most
likely to be found in typical organic molecules
are included in Table 2.2. The numbers (Pauling
electronegativity values) are on an arbitrary scale
from Li = 1 to F = 4.
From the sequence shown, it is readily seen
that hydrogen and carbon are among the least
electronegative atoms we are likely to encounter in
organic molecules. The relatively small difference
in electronegativities between hydrogen and carbon
also means there is not going to be much polarity
associated with a C–H bond. Most atoms other than
hydrogen and carbon when bonded to carbon are
going to be electron rich; therefore, bonds may
Table 2.2 Pauling electronegativity values
H
2.1
Li
Be
B
C
N
O
F
1.0
1.6
2.0
2.5
3.0
3.5
4.0
Na
Mg
Al
Si
P
S
Cl
0.9
1.2
1.5
1.8
2.1
2.5
3.0
K
B r
0.8
2.8
I
2.5
display considerable polarity. This polarity helps us
to predict chemical behaviour, and it is crucial to our
prediction of chemical mechanisms.
C O
C N
d+ d−
d+ d−
polarity in
and
bonds
C O
C N
We must also modify our thinking of bonding
as being simply ionic (where there is transfer of
electrons between atoms) or covalent (where there
is equal sharing of electrons). These represent two
extremes, but bond polarity now provides a middle
ground where there is sharing of electrons, but an
unequal sharing.
Bond polarity in a molecule can often be measured
by a dipole moment, expressed in Debye units
(D). However, the physical measurement provides
only the overall dipole moment, i.e. the sum of the
individual dipoles. A molecule might possess bond
polarity without displaying an overall dipole if two
or more polar bonds are aligned so that they cancel
each other out. The C–Cl bond is polar, but although
chloroform (CHCl 3 ) has a dipole moment (1.02 D),
carbon tetrachloride (CCl 4 ) has no overall dipole.
Because of the tetrahedral orientation of the dipoles
in carbon tetrachloride, the vector sum is zero.
H
Cl
C
Cl
Cl
overall dipole
Cl
Cl
C
Cl
Cl
tetrahedral orientation of
dipoles means vector sum
is zero
Polarization in one bond can also influence the
polarity of an adjacent bond. Thus, in ethyl chloride,
the polarity of the C–Cl bond makes the carbon more
positive (δ+); consequently, electrons in the C–C
bond are drawn towards this partial positive charge.
The terminal carbon thus also experiences a partial
positive charge, somewhat smaller than δ+ and so
depicted as δδ+.