Download Molecular Orbitals and Hybridisation

Organic Chemistry
Molecular Orbitals
and Hybridisation
Molecular orbitals
Orbitals can be used to explain bonding between
atoms. Atomic orbitals are the volume of space that
the electrons of an atom are likely to be found in.
H
H
1s atomic orbitals of hydrogen
The atomic orbitals containing the valence electrons
(outer electrons) are the ones that are important to us.
When atomic orbitals overlap, they combine to form
molecular orbitals.
In the case of hydrogen, the overlap of two 1s atomic
orbitals results in the formation of a σ (sigma)
molecular orbital.
H
H
H
H
σ molecular
1s atomic
orbitalsorbital
of hydrogen
σ
bonds
are covalent
atoms
This
molecular
orbital bonds
is moreformed
stablebetween
than each
of the
when
end-on
overlap
of orbitals
occurs. to the shape
separate
atomic
orbitals
and contributes
of the molecule.
Increasing energy
Molecular
The
molecular
orbitals
orbital
encompass
formed isthe
a lower
wholeenergy
molecule
arrangement
and
are not simply
than the
found
separate
between
atomic
atoms
orbitals.
inside a
molecule.
1s
1s
σ
H
H
H
H
Bonding continuum
The shape of the molecular orbital formed from
overlapping atomic orbitals will govern the type of
intermolecular bonding that is observed.
Non-polar (pure) covalent bonds
Non-polar covalent bonds or pure covalent bonds are
formed between two atoms of the same element, or
two atoms with a very low difference in
electronegativity.
The molecular orbital formed from overlapping atomic
orbitals is symmetrical around a mid-point where the
bonding electrons are most likely to be found.
Non-polar (pure) covalent bonds
Example: fluorine
The overlap of two 2p orbitals results in the formation
of a σ orbital.
F
+
F
2p atomic orbital
2p atomic orbital
σ molecular orbital
Non-polar (pure) covalent bonds
Example: fluorine
The overlap of two 2p orbitals results in the formation
of a σ orbital.
+
F
2p atomic orbital
F
F
2p atomic orbital
F
σ molecular orbital
In a fluorine molecule, or any non-polar covalent bond,
the σ bonding orbital is symmetrical.
Polar covalent bonds
Since oxygen
more difference
electronegative
thanthe
hydrogen,
When
there is is
a large
between
the molecular orbital
be asymmetrical,
with
electronegativities
of formed
the two will
elements
involved in the
the bonding
electrons
more likely
to will
be found
around
bond,
the bonding
molecular
orbital
be
the δ– oxygen atom.
asymmetrical.
δ–
Example: water
δ+
H
O
δ+
H
Ionic bonds
When ionic bonds form, there is extreme asymmetry
and the bonding molecular orbital is almost entirely
around one atom.
Hybridisation
In its ground state, an isolated atom of carbon has the
electron arrangement 1s2 2s2 2p2.
1s
H
C
H
H
H
2s
2p
Why then, if there are
only two unpaired
electrons, do carbon
atoms form four
covalent bonds?
Hybridisation
The shapes of the atomic orbitals involved cannot
explain the bonding observed in compounds such as
alkanes.
1s
2s
2p
z
z
y
z
1s orbital
2px orbital
y
y
z
2py orbital
2s orbital
y
z
2pz orbital
y
Increasing energy
Hybrid theory assumes that the 2s and 2p orbitals of
carbon atoms combine (or mix) to form four
degenerate orbitals (i.e. orbitals of equal energy)
2p
hybridised orbitals
2s
The hybrid orbitals formed
from one s orbital and
three p orbitals are called
sp3 orbitals.
an sp3 hybridised orbital
The sp3 orbitals formed are all half-filled, with the
electron far more likely to be found in the larger lobe.
Since electrons repel each
other, the four sp3 hybridised
orbitals surrounding a
central carbon atom result in
a familiar tetrahedral shape,
with a maximum possible
angle between each orbital
of 109.5°.
Alkanes
In methane, all four hybrid orbitals are used to form σ
bonds between the central carbon atom and hydrogen
atoms.
H
C
H
H
H
Alkanes
Carbon-to-carbon
single
bonds
in alkanes
result from
σ bonds are covalent
bonds
formed
by end-on
overlapping
sp3atomic
orbitals
forming
σ bonds.
overlap of two
orbitals
and
since σ bonds must
lie along the line joining both atoms, there will be free
rotation around these orbitals.
H
H
C
C
H
H
σ bond
H
H
Alkenes
How can we explain the existence of double bonds as
observed in alkenes?
H
H
H
C
C
H
H
H
H
H
C
H
C
H
As with alkanes, bonding in alkenes is due to
hybridisation.
As
alkanes,
electron of
from
thehybrid
2s shell
is
Thiswith
results
in theanformation
three
orbitals,
promoted
to the empty
2p orbital.2p orbital.
with one remaining
unhybridised
Increasing energy
single unhybridised
2p orbital
2p
hybridised orbitals
2s
The hybrid orbitals formed
from one s orbital and two p
orbitals are called sp2
orbitals.
The three sp2 orbitals repel each other, resulting in a
bond angle of 120° between them.
The hybrid orbitals are responsible for overlapping to
form σ bonds joining their central carbon atoms to
both carbon and hydrogen.
sp2 orbitals
The unhybridised p orbitals are perpendicular to the
plane of the molecule.
unhybridised
2p orbitals
H
H
C
σ bonds
H
C
σ bond
σ bonds
H
The p orbitals of the carbon atoms are parallel and
close enough to overlap sideways.
This sideways overlap between the 2p orbitals
produces a new molecular orbital between the two
carbon atoms.
A π new
bondorbital
is a covalent
formed
by
This
is calledbond
a pi (π)
orbital
theorsideways
overlap of
two
parallel
more commonly
aπ
bond.
atomic orbitals.
σ and π bonds
Looking at information comparing σ and π bonds, we
can see that double bonds are stronger than single
bonds, but not twice as strong. This is because the
sideways overlap (π bond) is weaker than the end-on
overlap (σ bond).
Bond type
Bonding
orbitals
present
Bond length
Mean
bond enthalpy
C
C
1σ
154 pm
370 kJ mol–1
C
C
1σ+1π
134 pm
602 kJ mol–1
C
C
1σ+2π
121 pm
835 kJ mol–1
Aromatic hydrocarbons
Aromatic compounds differ to other hydrocarbons as
they contain delocalised electrons.
Example: benzene (C6H6)
Chemists initially represented
H
a molecule of benzene as
H
H
C
shown here.
C
C
C
C
H
C
H
H
However, contrary to what
might be expected from this
structure, benzene is a very
stable, saturated structure
that does not undergo
addition reactions.
This model does not explain why all the bonds in
benzene can be observed to be the same length, not
three longer single bonds and three shorter double
bonds.
H
H
H
H
C
C
C
C
C
C
H
H
In benzene, each carbon
atom has used three of its
four valence electrons to form
σ bonds. The fourth electron
of each carbon atom is
delocalised over the entire
ring, not involved in π
bonding.
The σ bonding can be described as existing between
six sp2 hybridised orbitals.
There are six C–C σ bonds in the molecule and so
each carbon atom has two σ bonds to adjacent
carbon atoms.
Every carbon atom also has a σ bond to a hydrogen.
This results in a planar molecule with the unfilled 2p
orbital of each carbon atom above and below the
plane of the molecule.
H
H
H
C
C
C
C
H
C
C
H
H
These 2p orbitals all combine to form a set of
delocalised π molecular orbitals above and below the
plane of the molecule.
H
H
H
C
C
C
C
H
C
C
H
The structure of benzene is drawn as shown
to represent the delocalised electron clouds.
H
A substituted benzene ring is called a phenyl group
(C6H5) and can be represented:
R
Many medicines, antiseptics, drugs and other useful
products contain aromatic rings.
CH3
OH
NO2
O2N
NO2
Trinitrotoluene
(TNT)
Cl
Cl
Cl
Trichlorophenol
(TCP)