Quantum mechanics has given us the means determine the electron structure of isolated atoms. For instance, the electron structure of an isolated carbon atom can be written as: Energy 2p 2s 1s 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 form 4 bonds. Hybridization adjusts 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 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 Energy Energy 2p 2sp3 (hybrid orbital) hybridization 2s 1s 1s 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. Energy 2sp3 1s 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.

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). Energy Energy 2p 2sp3 2s 1s 1s In the hybrid orbital, there are three lone electrons to bond, and a lone pair . . . . . . . just like the Lewis structure.

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. Energy We can then reassign the five valence electrons to the new hybrid orbitals using Hund’ Rule. This creates 5 half filled orbitals capable of overlapping with the F atoms to form PF5 4s 3p 3s hybridization In the hybrid orbital, there are five lone electrons to bond. . . . . . . just like the Lewis structure. Energy 3d 4s 3sp3d (hybrid orbital)

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. domain = a bond set or a lone pair The domain of the molecule 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 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 one s orbital and three p orbitals.

Now see if you can use what you have learned to predict the hybridization of some other compounds. Let’s start with carbon dioxide Formula: CO2 Lewis Diagram: domain number: two domain geometry linear sp hybridization Hybridization Remember double or triple bonds count as 1 domain in VSEPR theory.

Now try the sulfate ion Formula: SO42- Lewis Diagram: domain number: four domain geometry : tetrahedral sp3 hybridization Hybridization

Now try the carbonate ion 2- Now try the carbonate ion 2- Formula: CO32- Lewis Diagram: 2- domain number: three domain name: trigonal planar sp2 hybridization Hybridization RESONANCE STRUCTURES – SWEET!!!!

Now try the sulfur hexafluoride Formula: SF6 Lewis Diagram: domain number: six domain shape: octohedral sp3d2 hybridization Hybridization