LOCALIZED AND DELOCALIZED BONDS
Atomic or hybrid orbitals overlap to form covalent bonds where the electrons are either localized or
delocalized.
"When the electrons forming a bond spend most of their time in the space between the two bonded
atoms, they are called localized electrons and such a bond is called localized bond, e.g., all -bonds
are localized bonds." On the other hand, "when the electrons are moving in and out of the space between the two bonded atoms, they are called delocalized electrons and the bond formed by them is called delocalized bond."
Electrons forming -bonds may be either localized or delocalized, e.g., -bond of ethylene is
localized because the electrons forming the -bond in ethylene are confined to the space between two
carbon atoms in such a way that these electrons are distributed equally in the space above and below
the plane of C—C, -bond (Fig. A). However, the two -bonds of acetylene are delocalized
because the electrons forming the -bond in the plane of paper do not remain confined to the space
above and below the plane of C—C, -bond and similarly the electrons of other -bond formed in
perpendicular direction, do not remain confined to that space. Actually, all these electrons merge together to form a cylindrical electron cloud around C—C -bond as in fig.
Other examples of delocalized bonds are 1, 3-butadiene and benzene where two or more than two -bonds are in conjugation. The electrons of one -bond are delocalized into the space of other -bond and vice-versa. This delocalization occurs through the overlap of unhybrid p orbitals present on each sp2 carbon. Thus;
(a) 1, 3-Butadiene: All the four carbon atoms of 1, 3-butadiene lie in the same plane due to which all the p-orbitals at four carbon atoms overlap with each other and the -electrons can move to a limited extent over all the four carbon atoms i.e., -electrons of 1, 3-butadiene are delocalized as shown in next Fig. (B). Hence, the -bonds of butadiene are delocalized bonds.
In contrast, the -electrons and hence the -bonds are localised in isolated dienes such as 1,4-pentadiene (next Fig. A), where each pair of -electrons is confined to the space between two carbon atoms.
(b) Benzene: In case of benzene there are six sp2-hybridized carbon atoms and each sp2-carbon
has one unhybridized p-orbital containing one electron. These p orbitals are so close to each other that they can overlap sideways to form a -bond. There are two modes for overlap of adjacent p-orbitals as shown in next Fig. ( A and B).
Actually, each p orbital overlaps equally well with the p orbitals on adjacent two carbon atoms on both sides to form a doughnut shaped -electron cloud above and below the plane of carbon and hydrogen atoms (see next Fig. ) i.e., these three -bonds of benzene are delocalized.
Armit and Robinson Structure
As the three -bonds of benzene are completely delocalized, it is not proper to represent benzene with a hexagonal ring with three double bonds at alternate positions since the position of -bonds is not fixed. Therefore, benzene is written as shown in last Fig. This representation was given by Armit
and Robinson.
Atomic or hybrid orbitals overlap to form covalent bonds where the electrons are either localized or
delocalized.
"When the electrons forming a bond spend most of their time in the space between the two bonded
atoms, they are called localized electrons and such a bond is called localized bond, e.g., all -bonds
are localized bonds." On the other hand, "when the electrons are moving in and out of the space between the two bonded atoms, they are called delocalized electrons and the bond formed by them is called delocalized bond."
Electrons forming -bonds may be either localized or delocalized, e.g., -bond of ethylene is
localized because the electrons forming the -bond in ethylene are confined to the space between two
carbon atoms in such a way that these electrons are distributed equally in the space above and below
the plane of C—C, -bond (Fig. A). However, the two -bonds of acetylene are delocalized
because the electrons forming the -bond in the plane of paper do not remain confined to the space
above and below the plane of C—C, -bond and similarly the electrons of other -bond formed in
perpendicular direction, do not remain confined to that space. Actually, all these electrons merge together to form a cylindrical electron cloud around C—C -bond as in fig.
(a) 1, 3-Butadiene: All the four carbon atoms of 1, 3-butadiene lie in the same plane due to which all the p-orbitals at four carbon atoms overlap with each other and the -electrons can move to a limited extent over all the four carbon atoms i.e., -electrons of 1, 3-butadiene are delocalized as shown in next Fig. (B). Hence, the -bonds of butadiene are delocalized bonds.
In contrast, the -electrons and hence the -bonds are localised in isolated dienes such as 1,4-pentadiene (next Fig. A), where each pair of -electrons is confined to the space between two carbon atoms.
has one unhybridized p-orbital containing one electron. These p orbitals are so close to each other that they can overlap sideways to form a -bond. There are two modes for overlap of adjacent p-orbitals as shown in next Fig. ( A and B).
As the three -bonds of benzene are completely delocalized, it is not proper to represent benzene with a hexagonal ring with three double bonds at alternate positions since the position of -bonds is not fixed. Therefore, benzene is written as shown in last Fig. This representation was given by Armit
and Robinson.
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