Hybridization as a Way of Explaining VSEPR Theory

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 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.

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.

Isolated Carbon atom Hybridized Carbon atom in a compound

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 half-filled orbitals capable of forming 4 bonds.

To summarize, we have seen that the atomic orbitals found in isolated atoms undergo a change (hybridization) when they are surrounded by other atoms in a compound.

These changes in the orbitals allow scientists to explain the bonding of various atoms in nature.

In the preceding example we were able to show why Carbon atoms frequently combine with other atoms to form compounds where the Carbon atom has 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 required the ammonia molecule to have a tetrahedral arrangement of its electron clouds. That would also yield 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.

hybridization

Isolated N atom N atom in ammonia

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.

hybridization

We can take 5 atomic orbitals from the P atom and “cross” them to form 5 equal hybrid orbitals.

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

(keep in mind that hybrid orbitals are always formed from the atomic orbitals in the valence shell)

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.

# of electron clouds

electron cloud geometry

type of hybridization (number)

atomic orbitals (formed from)

2 linear 2 orbitals (called sp hybrids) 1 s & 1 p3 trigonal planar 3 orbitals (called sp2 hybrids) 1 s & 2 p4 tetrahedral 4 orbitals (called sp3 hybrids) 1 s & 3 p5 trigonal bipyramidal 5 orbitals (called sp3d1 hybrids 1 s & 3 p & 1 d6 octohedral 6 orbitals (called sp3d2 hybrids) 1 s & 3 p & 2 d

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

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: four

Electron cloud geometry (angle between clouds):

Tetrahedral (<109o)

Hybridization sp3 hybridization

Now try carbon dioxide

Lewis Diagram:

# of electron clouds: Two

Linear (180o)

Hybridization sp hybridization

Remember double or triple bonds count as 1 electron cloud in VSEPR theory.

Electron cloud geometry (angle between clouds):

Now try the sulfate ion (SO4-2)

Lewis Diagram:

# of electron clouds: Four

tetrahedral (109o)

Hybridization sp3 hybridization

Electron cloud geometry (angle between clouds):

Now try the carbonate ion (CO3-2)

Lewis Diagram:

# of electron clouds: Three

Trigonal planar (120o)

Hybridization sp2 hybridization

Remember double or triple bonds count as 1 electron cloud in VSEPR theory.

Electron cloud geometry (angle between clouds):

Now try the sulfur hexaflouride (SF6)

Lewis Diagram:

# of electron clouds: Six

Octahedral (90o)

Hybridization sp3d2 hybridization

Electron cloud geometry (angle between clouds):