Hybridization

The content that follows is the substance of General Chemistry Lecture 35. In this lecture we Introduce the concepts of valence bonding and hybridization.

Valence Bond Theory

The Valence Bond Theory is the first of two theories that is used to describe how atoms form bonds in molecules. In this theory we are strictly talking about covalent bonds.

bond

According to the theory, covalent (shared electron ) bonds form between the electrons in the valence orbitals of an atom by overlapping those orbitals with the valence orbitals of another atom.

Bond Formation

When the bonds form, it increases the probability of finding the electrons in the space between the two nuclei.

There are two different types of overlaps that occur: Sigma (σ) and Pi (π)

Sigma (σ) Bonds form between the two nuclei as shown above with the majority of the electron density forming in a straight line between the two nuclei. I often refer to this as a "head-to-head" bond.

Pi (π) Bonds form when two un-hybridized p-orbitals overlap. This is what I call a "side-by-side" bond.

Pi (π) Bond

In order to overlap, the orbitals must match each other in energy. The process by which all of the bonding orbitals become the same in energy and bond length is called hybridization.

Hybridization

Let's start this discussion by talking about why we need the energy of the orbitals to be the same to overlap properly.

Let's look at the bonds in Methane, CH₄

The Carbon in methane has the electron configuration of 1s²2s²2p². According to Valence Bond Theory, the electrons found in the outermost (valence) shell are the ones we will use for bonding overlaps. This will be the 2s and 2p electrons for carbon.

As you know, p electrons are of higher energy than s electrons. This means that the two p electrons will make shorter, stronger bonds than the two s electrons right? But this is not what we see. We see a methane with four equal length and strength bonds. So how do we explain this? Simple: Hybridization

Carbon Hybridization

One of the s orbital electrons is promoted to the open p orbital slot in the carbon electron configuration and then all four of the orbitals become "hybridized" to a uniform energy level as 1s + 3p = 4 sp³ hybrid orbitals.

Hybridization process

Identifying Hybridization in Molecules

Figuring out what the hybridization is in a molecule seems like it would be a difficult process but in actuality is quite simple. Because hybridiztion is used to make atomic overlaps, knowledge of the number and types of overlaps an atom makes allows us to determine the degree of hybridization it has. In other words, you only have to count the number of bonds or lone pairs of electrons around a central atom to determine its hybridization.

The following rules give the hybridization of the central atom:
1 bond to another atom or lone pair = s (not really hybridized)
2 bonds to another atom or lone pairs = sp
3 bonds to another atom or lone pairs = sp²
4 bonds to another atom or lone pairs = sp³
5 bonds to another atom or lone pairs = sp³d
6 bonds to another atom or lone pairs = sp³d²

This Video Explains it further:

What are the hybridizations for each of the central atoms in the following molecule?

As you can see from the example above, assigning the hybridization to each central atom is easy as long as you can count to 6. What is really cool about the hybridization is that each hybridization corresponds to an electron pair geometry. So if you know the hybridization of an atom you automatically know its EPG.

Hybridization and Electron Pair Geometry

For s and sp hybridized central atoms the only possible molecular geometry is linear, correspondingly the only possible shape is also linear:

For sp2 hybridized central atoms the only possible molecular geometry is trigonal planar. If all the bonds are in place the shape is also trigonal planar. If there are only two bonds and one lone pair of electrons holding the place where a bond would be then the shape becomes bent.

For sp3 hybridized central atoms the only possible molecular geometry is tetrahedral. If all the bonds are in place the shape is also tetrahedral. If there are only three bonds and one lone pair of electrons holding the place where a bond would be then the shape becomes trigonal pyramidal, 2 bonds and 2 lone pairs the shape is bent.

For sp3d hybridized central atoms the only possible molecular geometry is trigonal bipyramidal. If all the bonds are in place the shape is also trigonal bipyramidal. If there are only four bonds and one lone pair of electrons holding the place where a bond would be then the shape becomes see-saw, 3 bonds and 2 lone pairs the shape is T-shaped, any fewer bonds the shape is then linear.

For sp³d² hybridized central atoms the only possible molecular geometry is Octahedral. If all the bonds are in place the shape is also Octahedral. If there are only five bonds and one lone pair of electrons holding the place where a bond would be then the shape becomes Square pyramid, 4 bonds and 2 lone pairs the shape is square planar, 3 bonds and 3 lone pairs the shape is T-shaped. Any fewer bonds the shape is then linear: