CH 106 - Lesson 2
Home Up Hybridization Covalent Bonding Sigma Bonds Types of Hybridization Double Bonds Pi Bonds Triple Bonds Comparing Hybrid Types Valence



To get into the concept of hybridization, I want to start with the tetrahedral arrangement of electrons around carbon atoms (which was introduced in Lesson 9 of CH-104). Then we will move on to the historical development of the tetrahedron concept and the historical development of the orbital concept. Finally, we will look at how the concept of hybrid orbitals connects these developments. 

Tetrahedral Methane

Let's begin by going back to the nice simple methane (CH4) molecule and talk about why the carbon bonds come out from the carbon atom in the directions that they do.

Model of methane.[62mod01.JPG]

The shape of CH4 is tetrahedral. We talked about it before a long time ago when we dealt with the shapes of very simple molecules. We related that shape to the electron dot diagram of CH4 in which there are eight dots around the carbon, arranged as four pairs of electrons, all of which are bonded to hydrogen atoms.


H : C : H


We interpret this by saying that if there are four groups of electrons around the center atom which are all bonded to other atoms, then the shape of the molecule is tetrahedral because that shape allows each electron group to get as far away from the other groups as possible. Perhaps I should say that shape results from each electron group getting as far away from the other groups as possible.

Well, this is a nice geometric and electrostatic argument, but it is not really the way chemists figured out that methane molecules, CH4, had a tetrahedral shape. That was figured out even before much, if anything, was known about electrons.

Tetrahedron History

Chemists in the 1800's determined that when one of the H atoms in methane was replaced by a Cl atom, it made no difference which H atom was replaced. They were all the same. Further, when a second H atom was replaced, it made no difference which one. Each of the three H atoms remaining after the first replacement are equivalent to one another. They have the same orientation with respect to the carbon and chlorine atoms. The only structure for methane and its derivatives that would allow this to happen is a tetrahedral structure with the C atom in the center and an H or other atom at each of the points of the tetrahedron.

You might want to do this yourself with a model when you are in the lab. Take off one of the hydrogens and replace it with a chlorine. You can use a green "atom" to represent chlorine. Then look at the model to verify for yourself that it really does not make a difference which hydrogen atom was replaced. Each one of those hydrogen atoms is the same with respect to all the others. Model of chloromethane.[62mod03.JPG]
Next, replace another hydrogen with a chlorine. Again, note that it really does not make any difference which one you selected. The orientation of the two chlorine atoms is going to be the same with respect to one another no matter which of the hydrogens you take off next. So, do this when you are in the lab. Model of dichloromethane.[62mod04.JPG]
Model of dichloromethane, different angle.[62mod05.JPG]


Flat Molecule Options

Now there is no way to explain that kind of property using flat models or any other shape.

For example, let's also consider a flat molecule with a carbon in the middle and hydrogen atoms going up and down and left and right as shown here at the bottom of this picture. (This is also shown in Example 1 in your workbook.) Tetrahedral and flat models of methane.[62mod06.JPG]
If you take one hydrogen atom off and replace it with a chlorine atom, you now have two different kinds of hydrogen as far as orientation to the chlorine (adjacent and opposite). Tetrahedral and flat models of chloromethane. [62mod07.JPG]
When you replace a second hydrogen atom, it makes a difference whether it is next to the chlorine atom or opposite the chlorine atom with the carbon atom between them. Two chlorine atoms opposite one another is a different structural arrangement than two chlorine atoms next to one another. Tetrahedral and flat models of dichloromethane. [62mod08.JPG]

So, a flat arrangement for CH4 does not explain the property of being able to replace any two hydrogens with chlorine without it making any difference which ones are replaced.

Orbital History

In the late 1800's and early 1900's more and more was learned about electrons. You should remember that the number and arrangement of electrons in the atoms of the different elements was related to the shape of the periodic table and the location of the elements on it.

We talked about electron configurations and determined that the electron configuration for carbon is 1s22s22p2. Section of periodic table with electron configuration for carbon. [62pt01.JPG]
This created a dilemma for theoretical chemists. How can the 1s22s22p2 arrangement of electrons in the orbitals of a carbon atom be related to the arrangement of electrons in CH4 to give a tetrahedral shape? Section of periodic table with electron configuration for carbon and model of methane. [62pt02.JPG]

The answer quite simply is that it cannot, at least not directly. Nevertheless, they persisted and came up with a relationship or process that you need to know about. It is called hybridization.

First, let me set the stage. Just as you can alter your clothing (e.g. zip up a jacket and raise its hood, etc.) to respond to different situations, so an atom can alter the arrangement of its electrons to fit the conditions in which it finds itself. You know that electrons take up space and that the space they take up is referred to as an orbital. There are many kinds of orbitals. In the past we talked about s, p, d and f orbitals. These are all atomic orbitals. They represent the ways that electrons can arrange themselves in isolated, individual atoms. When atoms bond to one another, the electrons have to change their arrangement in order to accommodate influence of other atoms. The space taken up by bonded electrons is called a bonding orbital.

To review, atomic orbitals are used to explain the spectra of individual atoms and the shape of the periodic table. Bonding orbitals are used to explain the shape and properties of molecules consisting of atoms sharing electrons.

Hybrid Orbitals

Enter hybridization and hybrid orbitals. Hybridization is the name we use to describe the process of change from atomic orbitals to bonding orbitals. We refer to the orbitals that have been changed as hybrid orbitals.

Here is the idea: when a carbon atom bonds to other atoms, the four orbitals in the second shell are somehow mixed together and rearranged to give four new orbitals. These four new orbitals are called hybrid orbitals. Each of the four hybrid orbitals is equivalent to the others and each contains one electron.

Naming hybrid orbitals has a tricky angle to it. If you think of hybridization having a "before" mode and an "after" mode, the hybrid orbitals are the "after" mode, but their name comes from the "before" mode. Because these hybrid orbitals resulted from the combination of an s orbital and three p orbitals, they are called sp3 hybrid orbitals or simply sp3 orbitals. A local parallel in this type of naming comes to mind. The "Berryhill" and "Holly Farm" shopping centers are named for what used to be in those locations. Similarly, hybrid orbitals are named for the atomic orbitals that the electrons used to occupy.


Here is one way of picturing hybridization in terms of what we have done before. (It is also shown in Example 2 in your workbook.) From the electron configuration of carbon we know that carbon has four valence electrons, two in the 2s orbital and two in the 2p orbitals. The valence electrons can be represented as dots in the Lewis diagram. When we draw the electron dot diagram for carbon, the four dots representing the four valence electrons are usually drawn separately and spread around the C. We didn't talk about it this way at the time, but the process of separating and rearranging the electrons is what we can now call hybridization. Then, those rearranged electrons bond to other atoms.
1s22s22p2   or
: C


H : C : H



Now let's look at hybridization in terms of orbitals. The s and three p orbitals of the carbon's second energy level combine and rearrange to form four new hybrid orbitals. This fuzzzy drawing shows the orbitals which point out from the center of the atom in the same direction as the corners of a tetrahedron. Diagram of sp3 orbitals. [62orb00.JPG]
In this diagram a little bit more detail about the shape of those hybrid orbitals is shown. Just one orbital is shown here. Like a p orbital, it has two lobes; but unlike a p orbital, one lobe is much larger than the other and the electron will spend almost all of its time in the larger portion of the orbital. Diagram of single sp3 orbital. [62orb01.JPG]

In fact, the smaller part of the orbital is often left out of drawings. You can see that in the previous diagram only the larger lobes (where the electrons spend most of their time) are shown.

Hopefully, this page has given you an awareness of the hybridization process. The following pages of this section will show how those hybrid orbitals are involved in the bonding in organic compounds.

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E-mail instructor: Eden Francis

Clackamas Community College
2001, 2003 Clackamas Community College, Hal Bender