BOND POLARITY
35
π bond
π bond
formation of one
and
two
σ bonds, plus one
π bond in formaldehyde
C
O
H
H
π bond
π bond
C
N
H
H
π bond
π bond
π bond
π bond
formation of one
, two
H
C
O
H
H
side view
top view
C
N
H
H
top view
side view
H
C O
C N
C H
C O
C H
C O
C O
C N
C N
C N
N H
π bond in methanimine
and one
σ bond, plus one
Figure 2.22 Bonding in formaldehyde and methanimine
π bonds
π bonds
formation of one
and one
σ bonds, plus two
π
bonds in hydrogen cyanide
C
N
π bonds
π bonds
C
N
side view
H
H
C N
C N
C H
C N
C N
Figure 2.23 Bonding in hydrogen cyanide
oxygen carries two lone pairs in sp
2 orbitals, whereas
nitrogen carries one. Thus, the main difference from
the alkene structure, apart from the atoms involved,
is that lone pairs in atomic orbitals replace one or
more of the σ molecular orbitals that constituted the
C–H bonds. The atoms around the double bond are
in a planar array, just as in an alkene.
Triple bonds are also encountered in cyanides/nitriles. We can compare these with alkynes in much
the same way (see Section 2.6.2). With sp-hybridized
nitrogen, we can form one C–N σ bond and two
C–N π bonds, leaving a lone pair of electrons on
nitrogen in an sp atomic orbital (Figure 2.23). The
cyanide/nitrile system is linear, just like an alkyne.
2.7 Bond polarity
The nucleus of each atom has a certain ability to
attract electrons. This is termed its electronegativity.
This means that, when it is bonded to another atom,
the bonding electrons are not shared equally between