Preferred Label : σ, π (sigma, pi);
IUPAC definition : The terms are symmetry designations, π molecular orbitals being antisymmetric with
respect to a defining plane containing at least one atom (e.g. the molecular plane
of ethene) and σ molecular orbitals symmetric with respect to the same plane. In practice
the terms are used both in this rigorous sense (for orbitals encompassing the entire
molecule) and also for localized two-centre orbitals or bonds, and it is necessary
to make a clear distinction between the two usages. In the case of two-centre bonds,
a π-bond has a nodal plane that includes the internuclear bond axis, whereas a σ-bond
has no such nodal plane. (A δ-bond in organometallic or inorganic molecular species
has two nodes.) Radicals are classified by analogy into σ- and π-radicals. Such two-centre
orbitals may take part in molecular orbitals of σ- or π-symmetry. For example, the
methyl group in propene contains three C–H bonds, each of which is of local σ-symmetry
(i.e. without a nodal plane including the internuclear axis), but these three 'σ-bonds'
can in turn be combined to form a set of group orbitals one of which has π-symmetry
with respect to the principal molecular plane and can accordingly interact with the
two-centre orbital of π-symmetry (π-bond) of the double-bonded carbon atoms, to form
a molecular orbital of π-symmetry. Such an interaction between the CH sub 3 /sub
group and the double bond is an example of what is called hyperconjugation. This cannot
rigorously be described as 'σ–π conjugation' since σ and π here refer to different
defining planes, and interaction between orbitals of different symmetries (with respect
to the same defining plane) is forbidden.;
Origin ID : S05434;
See also
The terms are symmetry designations, π molecular orbitals being antisymmetric with
respect to a defining plane containing at least one atom (e.g. the molecular plane
of ethene) and σ molecular orbitals symmetric with respect to the same plane. In practice
the terms are used both in this rigorous sense (for orbitals encompassing the entire
molecule) and also for localized two-centre orbitals or bonds, and it is necessary
to make a clear distinction between the two usages. In the case of two-centre bonds,
a π-bond has a nodal plane that includes the internuclear bond axis, whereas a σ-bond
has no such nodal plane. (A δ-bond in organometallic or inorganic molecular species
has two nodes.) Radicals are classified by analogy into σ- and π-radicals. Such two-centre
orbitals may take part in molecular orbitals of σ- or π-symmetry. For example, the
methyl group in propene contains three C–H bonds, each of which is of local σ-symmetry
(i.e. without a nodal plane including the internuclear axis), but these three 'σ-bonds'
can in turn be combined to form a set of group orbitals one of which has π-symmetry
with respect to the principal molecular plane and can accordingly interact with the
two-centre orbital of π-symmetry (π-bond) of the double-bonded carbon atoms, to form
a molecular orbital of π-symmetry. Such an interaction between the CH sub 3 /sub
group and the double bond is an example of what is called hyperconjugation. This cannot
rigorously be described as 'σ–π conjugation' since σ and π here refer to different
defining planes, and interaction between orbitals of different symmetries (with respect
to the same defining plane) is forbidden.