Transcript Simple bonding models and tools for predicting shapes

Hybridization
A Way of Explaining
VSEPR Theory
Covalent bonding

Modern methods for describing bonding
make use of quantum mechanical
methods and describe the electrons in
molecules in terms of molecular orbitals.

However, simpler earlier theories such as
that due to Lewis can be useful in some
cases and they are a lot less complex.
Quantum mechanics has given us the means
(Schrodinger’s Wave Equation) to be able to
determine the electron structure of isolated atoms.
For instance, the electron structure of an
isolated Carbon atom can be written as:
Based on this model, we would predict that
Carbon atoms should form two covalent
bonds (since covalent bonds involve the
overlap of half-filled orbitals).
However, in nature we find that Carbon
atoms do not normally behave in this
manner. Instead Carbon atoms almost
always form 4 bonds.
Hybridization theory is an attempt by
chemists to adjust the concept of
electron structures of various atoms,
including Carbon atoms, to make them
consistent with the way they are
observed to bond in nature.
The basic idea of hybridization is indicated by
the name. A hybrid in biology is an offspring of
parents with different characteristics. For
instance, a mule is a hybrid of a a horse and a
donkey.
Hybridization in chemistry involves
combining atomic orbitals to form a new set
of “hybrid” orbitals. These new orbitals will
have some of the properties of the different
atomic orbitals which go into forming them.
Let’s take a Carbon atom to see how the formation
of the hybrid orbitals creates a new set of orbitals
with some properties of the atomic orbitals.
Hybridized Carbon atom in a
compound
Isolated Carbon atom
hybridization
Notice that the hybrid orbitals are “crosses” between the
low energy s orbital and the higher energy p orbitals.
Also notice that the hybrid orbitals all are equal in
energy.
After forming the hybrid orbitals, the electrons
must be distributed among the new orbitals.
Since these hybrid orbitals are equal in energy,
the electrons must distributed according to
Hund’s Rule.
In the case of a Carbon atom, the four
valence electrons are distributed among the
four hybrid orbitals. This produces 4 halffilled orbitals capable of forming 4 bonds.
However, we need to remember that Carbon atoms are
not the only atoms to undergo hybridization. Let’s
look at the electron configuration for an isolated
Nitrogen atom and see if hybridization can be used to
explain how it bonds to form ammonia (NH3).
At first glance it might appear that there is no need for
hybridization since the Nitrogen atom already seems to
have the ability to form 3 covalent bonds. However, there
are difficulties with using the atomic orbitals to explain
the bonding in ammonia.
If we assume that the Nitrogen bonds due to
the overlapping of its p orbitals then we
should find that the bond angle in ammonia
would be 90o since the p orbitals are located
on the x,y, and z axis.
However, this explanation breaks down when we
discover that the experimentally determined
bond angles in ammonia are approximately 107o
However, if we utilize VSPER Theory to predict
the shape, we get a more satisfying prediction. To
utilize VSPER Theory we must first determine
the Lewis Structure.
The four electron clouds in the structure would
have a tetrahedral arrangement with a
predicted bond angle of 109o, which is in close
agreement to the experimentally measured
bond angle.
Let’s assume that the nitrogen atom undergoes
hybridization in a manner similar to carbon.
N atom in ammonia
Isolated N atom
hybridization
This approach yields four equal orbitals, three
½ filled and one full, which is consistent with
the VSEPR prediction as well as the
experimentally determined bond angle.
The hydrogen atoms overlap on the three ½
filled hybrid orbitals and the other hybrid
orbital contains a non-bonding pair of
electrons.
Even the slight difference between the experimental
bond angle (107o) and the theoretical angle (109o)
can be explained by the fact that the repulsion of the
lone pair electrons is greater than for the bonding
electrons.
Hybridization Theory is also capable of explaining
molecules such as PF5 which contain expanded
octets.
If we determine the Lewis diagram for this
molecule, we find the following:
It is not possible to explain this structure without
hybridization theory!
However, by utilizing hybridization we can explain
how the phosphorus atom is able to form 5 bonds.
We can take 5 atomic orbitals from the P atom and
“cross” them to form 5 equal hybrid orbitals.
(keep in mind that hybrid orbitals are always formed
from the atomic orbitals in the valence shell)
hybridization
We can then reassign the five valence electrons
to the new hybrid orbitals using Hund’ Rule.
This creates 5 ½ filled orbitals capable of
overlapping with the F atoms to form PF5
Hybridization theory is a way of explaining the
shapes of molecules which are found in nature
(and predicted by the VSEPR theory.
The different shapes can be explained by the
different types of hybridization.
The shape of the molecule (as determined
by the VSEPR theory) determines the type
of hybridization which the central atom
must undergo.
The following chart shows the type of
hybridization which can be used to explain the
various shapes found in nature and predicted by
the VSEPR theory
# of electron
clouds
2
3
4
5
6
electron cloud
type of hybridization
geometry
(number)
linear
2 orbitals (called sp hybrids)
trigonal planar
3 orbitals (called sp2 hybrids)
tetrahedral
4 orbitals (called sp3 hybrids)
trigonal bipyramidal 5 orbitals (called sp3d1 hybrids
octohedral
6 orbitals (called sp3d2 hybrids)
atomic orbitals
(formed from)
1s&1p
1s&2p
1s&3p
1s&3p&1d
1s&3p&2d
Notice that the name of the hybrid orbitals is
determined by the atomic orbitals which were
combined to form them.
For instance: sp3 hybrids were formed from 1 s
orbital and 3 p orbitals.
Now see if you can use what you have learned
to predict the hybridization of some other
compounds. Let’s start with water.
Lewis Diagram:
# of electron clouds:
Hybridization
four
sp3 hybridization
Now try carbon dioxide
Lewis Diagram:
# of electron clouds:
Hybridization
Two
sp hybridization
Remember double or triple bonds count
as 1 electron cloud in VSEPR theory.
Now try the sulfate ion (SO4-2)
Lewis Diagram:
# of electron clouds:
Hybridization
Four
sp3 hybridization
Now try the carbonate ion (CO3-2)
Lewis Diagram:
# of electron clouds:
Hybridization
Three
sp2 hybridization
Remember double or triple bonds count as 1
electron cloud in VSEPR theory.
Now try the sulfur hexaflouride (SF6)
Lewis Diagram:
# of electron clouds:
Hybridization
Six
sp3d2 hybridization