Hybrid Orbitals explained - Valence Bond Theory | Crash Chemistry Academy

Hello and welcome to a video on

hybrid

orbitals

, often called

valence

bond

theory

. Developed in the 1930s by the great chemist Linus Pauling as a

bond

ing model to understand the three-dimensional placement of atoms in a molecule, and which is fundamental to our understanding of the properties that molecules have. In this video we will see methane, ethene, ethyne, ammonia and water as our

hybrid

ization models. There are only a small group of atoms in the second period for which the model actually works, but among them are carbon, nitrogen and oxygen, which make up the vast majority of molecules that exist on Earth.

Therefore, the model applies to a limited number of elements, but it applies to by far the majority of molecules. Let's take a look at the word hybrid: it is a mixture of two varieties. If you put a horse and a donkey together, you will get a mule, a mix or a hybrid of a horse and a donkey. Let's go to hybrid

orbitals

using carbon as a model. As a single atom not bonded to anything, carbon has two 1s electrons, two 2s electrons, and two 2p electrons. This electronic configuration is the energetic arrangement of the electrons of the carbons.

However, carbon rarely exists in nature as an individual atom, except momentarily while undergoing chemical reactions. Carbon exists with its

valence

electrons bonded to other atoms. When carbon is bonded to four other atoms, the four bonds of carbon are experimentally considered to be equivalent. And so, when carbon atoms are in a bonding situation, their own bonding electrons exist at equivalent energies, requiring them to hybridize with an energy intermediate between the 2s energy and the 2p energy. Or you can think of it as a combination or hybridization of the two energies, the s and p energies. And because the energies of these electrons have now changed, the shape of the orbitals they occupy is also different, which we will see in a moment, and they are called hybrid orbitals.

They are called 2sp3 hybrid orbitals. The naming often confuses students, so before we continue, let's take a look at where the name 2sp3 comes from: The 2 comes from the second major energy level in which the valence orbitals are found. The s comes from the 2s orbital that contributes to hybridization, and the p comes from the 2p orbitals that contribute to hybridization, and the 3 comes from the number of 2p orbitals used in hybridization. Once hybridized, the 2s and 2p orbitals no longer exist, so we have four 2sp3 hybrid orbitals. Four 2sp3 hybrid orbitals derived from combining the energies of one 2s orbital and three 2p orbitals, giving a total of four 2sp3 orbitals.

Before examining the shape of hybrid orbitals, it would be helpful to briefly review atomic orbitals. The 1s orbital is a sphere, the 2s orbital is a larger sphere surrounding the 1s, and here we will get rid of the 1s since we are only interested in the valence electrons. Each 2s orbital has a two-lobe shape that converges at the nucleus. So there are the three 2p orbitals. However, when hybridization occurs, the s and p orbitals cease to exist and the 2sp3 orbitals have a completely different shape. We can see that orbital hybridization explains the VSEPR placement of carbon's four valence electrons, since the four 2sp3 orbitals are equivalent, each 2sp3 orbital repelling the others with equal strength, resulting in identical bond angles.

The carbon atom only hybridizes when it is in a bonding situation. Here, four hydrogen atoms bond to carbon by overlapping their orbitals with the hybrid orbitals of carbon. So what would be the reason for this? If we step back and see that both carbon and hydrogen have unpaired electrons, the superposition allows the electrons to pair and therefore go to a lower potential energy. The illustration here contains the valence electrons of both carbon and hydrogen, and since everyone likes to visualize atoms as spheres, we can do the same: here is our carbon atom and here are the hydrogens, with the spheres superimposed, which indicating the overlapping orbitals that constitute the bond.

Bonds are more easily discernible in a ball and stick model, which also makes the bond angle more visible. Since the four sp3 orbitals are equivalent, each bonding orbital repels the others with the same strength, resulting in identical bond angles. Bonds in hybridization also have their own nomenclature. The overlapping orbitals are called sigma bonds and represent the single bond occupied by a single pair of electrons. What about double bonds? How does the hybridization model explain double bonds? We will use ethene, C2H4, to see what happens in a double bond. The single bonds we know are sigma bonds, and the double bond also has a sigma bond, but the second bond of a double bond is a pi bond.

Let's see how hybridization and orbital overlap can explain a double bond. The two carbon atoms in ethene are equivalent, so let's look at one of the carbon atoms first. A pi bond arises from the overlap of unhybridized p orbitals, so the atom hybridizes only three orbitals, leaving one unhybridized p orbital for the pi bond. The hybrid orbital is called 2sp2 and the superscript 2 indicates that only two 2p orbitals have contributed to the hybridization. The hybridized 2sp2 orbitals exist in a plane perpendicular to the unhybridized 2p orbital. Let's eliminate the 2p orbital for now to see it easier.

The hybridized 2sp2 orbitals are distributed at an angle of 120 degrees, meaning that they exist in a plane, and the plane is perpendicular to the unhybridized 2p orbital. So this is what both carbon atoms do when the bond occurs in ethene. Each carbon atom is sp2 hybridized. The sigma bond occurs with overlapping 2sp2 orbitals. What about the pi link? The second bond of the double bond. We previously said that it comes from the unhybridized p orbitals, which we see here from both carbon atoms. The upper and lower lobes of the 2p orbitals overlap above and below the sigma bond axis forming a single pi bond.

The space in which the now paired electrons move. The ball-and-stick model shows this double bond with two dashes. In short, sigma bonds occur along the axis between nuclei. Pi bonding occurs above and below the Sigma axis where the p orbital lobes have overlapped. The ethene molecule also bonds to four hydrogen atoms overlapping with the other 2sp2 hybrid orbitals of both carbons, creating four more sigma bonds. In the ball and stick model we can easily see that each carbon has a trigonal planar geometry and therefore the entire molecule exists in a plane with the single pi bond above and below that plane.

Now let's see how hybridization can be a model for the triple bond using ethyne, C2H2. The carbon-hydrogen bonds are sigma bonds and the triple bond is one sigma bond and two pi bonds. Let's see how hybridization can adapt to this. Since pi bonds come from p orbitals, and we need two pi bonds, then two 2p orbitals must remain unhybridized, so the remaining single 2s orbital and a single 2p orbital will hybridize to two hybridized 2sp orbitals. And there are also the violet 2px and blue 2py orbitals. Here each green lobe is a single orbital, so each has one electron, while both purple lobes constitute the single 2px orbital with a single electron.

And both blue lobes constitute the only 2py orbital with a single electron. The other carbon in C2H2 also has the same triple bond, so it has the same hybridization. Let's see what happens during pairing. The sp orbitals of both carbons overlap, forming a sigma bond. The upper lobes of the blue 2p orbitals overlap, as do the two lower lobes, creating the first of the 2 pi bonds. Can you guess where the second of the two pi bonds comes from? Yes it's correct! It is the overlap of the 2p violet lobes. Let's get rid of the sigma bond for a moment to see something interesting.

Each pi bond lies in a separate plane and the two planes are perpendicular to each other, so the two pi bonds are perpendicular to each other. Finally, two hydrogens will overlap with the remaining sp hybrid orbitals, creating the C2H2 molecule. The overlap of the space-filling model reflects the overlapping orbitals, which is also represented by the ball-and-stick model. In the rest of the video we will look at the hybridization of nitrogen and oxygen using NH3, ammonia, as our model for the hybridization of nitrogen, and H2O, water, as our model for oxygen. In NH3, nitrogen has three sigma bonds and one lone pair, so how does this explain hybridization?

The hybridization is 2sp3 and nitrogen has 5 valence electrons, so one of the four hybrid orbitals of 2sp3 has a pair of electrons. The three sigma bonds come from sp3 orbitals with a single electron, so they can pair, so the remaining pair of electrons is a lone pair, a pair of unbonded electrons. As with sp3 hybridization in carbon, the hybrid orbitals of nitrogen extend tetrahedrally. And finally water. Here oxygen has two sigma bonds and two lone pairs. In water, oxygen also has 2sp3 hybridization, but with six valence electrons: two of the sp3 orbitals have paired electrons. You can probably guess that sp3 orbitals with a single electron will overlap with hydrogen, and the remaining two pairs are not bonded, they are lone pairs.

Again the hybrid orbitals of oxygen extend in a tetrahedral manner. That's all for hybridization, product of a mad scientist! SEE YOU!