Lesson Video: Hybridization Chemistry • Second Year of Secondary School

Video Transcript

In this video, we’ll learn about hybridization. We’ll learn how atomic orbitals hybridize and how to identify the type of hybridization in a molecule.

Before we dive into hybridization, let’s recall that atomic orbitals are mathematical functions that help us describe the wave-like properties of electrons and atoms. Because electrons have these wave-like properties, which are described by quantum mechanics, we can’t describe the electron as a point orbiting the nucleus. Instead, the electron is more like a cloud that spread out in space around the nucleus of the atom. The shape of this electron cloud is described by the mathematical function of the atomic orbital. Atomic orbitals have different shapes and sizes that are related to the energy of the electron.

s orbitals have a spherical shape, while p orbitals have more of a dumbbell shape and are oriented around one of the Cartesian coordinate axes. Orbitals come in other shapes and sizes as well. But the 1s, 2s, and 2p orbitals are the ones we’ll encounter in this video.

Recall the numbers in front of the orbital type indicate the energy level, and the letters tell us about the orbital shape. In an atom, electrons fill the orbitals in order of increasing energy. So the single electron in the hydrogen atom will occupy a 1s orbital. If this 1s orbital overlaps the 1s orbital of another hydrogen atom, a covalent bond is formed. Each atom participating in the bond contributes one electron to the bond. Specifically, this bond is called a sigma bond, often noted by the Greek lowercase letter 𝜎. 𝜎 bonds are formed when orbitals overlap end-to-end.

Now, let’s see if we can describe the bonding in a slightly more complicated molecule like methane using orbital overlap. Methane has four covalent carbon-hydrogen bonds. One of the electrons, in each of these bonds will come from the 1s orbital of hydrogen atom. The other electron in each bond must come from one of carbon’s valence electrons, which are located in the 2s and 2p atomic orbitals. So it seems some bonds in methane must be from the overlap of carbon’s 2s orbital and hydrogen’s 1s orbital and the rest from the overlap of carbon’s 2p orbital and hydrogen’s 1s orbital.

However, experimental evidence tells us that all four carbon-hydrogen bonds are identical. So this picture of orbital overlap we’ve come up with for methane must not be correct. Something else must be going on here to create these four identical carbon-hydrogen bonds.

It turns out that atomic orbitals get mixed together to create hybrid orbitals that have new and different shapes, and these new hybrid orbitals will then be used to form the bonds in a molecule. This mixing is not physical. Rather, it’s a mathematical mixing of the functions that describe the atomic orbitals.

Let’s see how orbital hybridization works by looking again at the carbon and methane. This is carbon’s valence electron configuration. Carbon has valence electrons in the 2s and 2p atomic orbitals. The first step in orbital hybridization is the promotion of an electron. For carbon, a 2s electron is promoted to an empty 2p orbital, which puts carbon in an excited state. Now, carbon has four unpaired electrons in the 2s and 2p orbitals. These four unpaired electrons in different orbitals are mixed together to create four identical orbitals of the same energy. These four identical orbitals are neither s orbitals or p orbitals. They are sp3 hybrid orbitals.

These sp3 hybrid orbitals have a similar oblong shape to the p orbitals, but they aren’t oriented around a Cartesian axis like the p orbitals are. Instead, the sp3 hybrid orbitals are oriented so that each lobe points towards the corner of a tetrahedron.

In methane, carbon’s sp3 hybrid orbitals can overlap the 1s orbitals of hydrogen, creating the four identical carbon-hydrogen bonds in methane. Because of the tetrahedron geometry of the sp3 hybrid orbitals, all of the carbon-hydrogen bond angles in methane are 109.5 degrees. This bond angle maximizes the distance between the electron pairs in the bonds, which minimizes the electrostatic repulsions between these electron groups.

Now, let’s take a look how orbitals hybridize in a different molecule, this time boron trifluoride. This is the ground state electron configuration for boron’s valence electrons. There are electrons in the 2s and one of the 2p atomic orbitals. Just as before, one of the electrons in the 2s orbital will become excited, which promotes it to a vacant 2p orbital.

Now, there are three unpaired electrons in the 2s and 2p orbitals. These will be hybridized to create three identical sp2 hybrid orbitals. The sp2 hybrid orbitals are oriented so that each lobe points towards one of the corners of a triangle. The sp2 hybrid orbitals can overlap with the 2p orbitals of fluorine that have an unpaired electron. This creates three identical boron-fluorine bonds with a 120-degree bond angle. And just as we saw in methane, this bond angle maximizes the distance between the bonding electrons, which minimizes the electrostatic repulsion between them.

Let’s take a look at one last example of orbital hybridization with the molecule beryllium hydride. We’ll again start off with the basic electron configuration for beryllium’s valence electrons. Both of beryllium’s valence electrons are in a 2s orbital. Because both of these electrons are paired, beryllium can’t form any bonds with hydrogen unless it undergoes hybridization. Once again, this process starts with the promotion of one of those 2s electrons to an empty 2p orbital, putting beryllium in an excited state. The unpaired electrons in the 2s and 2p atomic orbitals are hybridized, leaving us with two identical sp hybrid orbitals. These two sp hybrid orbitals are oriented with their lobes pointing in opposite directions along a line.

When the two sp hybrid orbitals overlap with hydrogen’s 1s orbital, the result is a linear molecule with 180-degree bond angles. This bond angle will place the bonding electrons on opposite sides of the molecule, which is the maximum distance they can be away from each other. This once again will minimize the repulsions between those electron groups.

Before we move on, it’s worth noting that this first step in the hybridization process where an electron is promoted requires energy. But the energy required to promote an electron is much less than the amount of energy that’s released from the formation of bonds. Sometimes, not all of the atomic orbitals are involved in the hybridization process, leaving some orbitals unhybridized. We see this in the molecule ethene. We know that ethene is a flat molecule with 120-degree bond angles, which suggests sp2 hybridization.

Let’s walk through the hybridization process for one of the carbons in ethene. We’ll start off as always with the valence electrons of the carbon in their basic state. One of the 2s electrons will be promoted, putting the carbon in an excited state. The unpaired electrons in the 2s and two of the 2p atomic orbitals will be hybridized to create three sp2 hybrid orbitals. But one of the 2p orbitals will remain unhybridized. The sp2 orbitals from each carbon will overlap with the 1s orbital of a hydrogen atom, creating the four carbon-hydrogen bonds in ethene. The sp2 orbitals can also overlap with each other, creating one of the carbon-carbon bonds.

The remaining carbon-carbon bond will be formed using the remaining unhybridized 2p orbitals. These orbitals overlap side-by-side, creating the second carbon-carbon bond. This bond looks very different from the other bonds we’ve seen in this molecule and in the other molecules in this video, for example, the carbon-hydrogen bonds and the carbon-carbon bonds. The rest of the bonds we’ve looked at have featured the end-to-end overlap of orbitals, which, as we said earlier, makes these 𝜎 bonds. But this bond is formed by the side-by-side overlap of adjacent p orbitals. This kind of bond is called a 𝜋 bond.

The electron density of a 𝜋 bond is split above and below the molecule with a notable plane or a region with no electron density in the plane of the carbon and hydrogen atoms. Whenever we have a molecule with a double bond, one bond will be a 𝜎 bond and the other will be a 𝜋 bond.

We can see something very similar in the molecule ethyne. This linear molecule suggests sp hybridization. When the atomic orbitals of a carbon and ethyne are hybridized, we end up with two electrons in sp hybrid orbitals and two electrons in unhybridized 2p orbitals. The sp hybrid orbitals overlap with hydrogen’s 1s orbitals to create the carbon-hydrogen bonds. And the sp hybrid orbitals from each carbon overlap to create one of the carbon-carbon bonds. All three of these bonds are 𝜎 bonds. The remaining two carbon-carbon bonds are 𝜋 bonds formed from the two unhybridized 2p orbitals. So triple bonds, like the one in ethyne, are formed from one 𝜎 and two 𝜋 bonds.

Now we’ve learned quite a bit about hybridization. So let’s try some problems.

During the formation of chemical bonds, hybridization can occur. Atomic orbitals merge together mathematically to form hybrid orbitals. What hybrid orbitals are formed when one s orbital and three p orbitals all hybridize? (A) Four sp2 orbitals, (B) two sp2 orbitals, (C) four sp3 orbitals, (D) one sp3 orbital, (E) three sp3 orbitals.

As the question states, hybridization is when the mathematical functions described by the atomic orbitals are mixed together to form hybrid orbitals. These hybrid orbitals have new and different shapes. Say we have electrons in s and p orbitals. The hybridization process begins with one of the electrons in an s orbital getting promoted to an empty p orbital. These four unpaired electrons, one in the s orbital and three in the p orbital, will merge together to create four identical sp3 hybrid orbitals.

We can start off with a different number of unpaired electrons to create different kinds of hybrid orbitals. Unpaired electrons in 1s orbital and 2p orbitals will hybridize to give us three sp2 orbitals. And unpaired electrons in 1s and 1p orbital hybridize to give us two sp orbitals. This question asked us which hybrid orbitals are formed, when 1s and 3p orbitals hybridize. As we just saw, this will give us four sp3 orbitals, which is answer choice (C).

A 𝜎 bond is created when specific combinations of atomic orbitals overlap. Which of the following occurrences does not result in the formation of a 𝜎 bond? (A) A hybrid sp orbital overlapping with another hybrid sp orbital, (B) an s orbital overlapping with a p orbital, (C) a p orbital overlapping with another p orbital side-by-side, (D) an s orbital overlapping with another s orbital, (E) a p orbital overlapping with another p orbital end-to-end.

Sigma bonds, often noted with the lowercase Greek letter, are formed by the end-to-end overlap of orbitals. It doesn’t matter which two orbitals are overlapping as long as they overlap end-to-end. Two s orbitals overlapping can create a 𝜎 bond or an s orbital can overlap with a p orbital. A 𝜎 bond results from the s orbital overlapping any of the hybrid orbitals. 𝜎 bonds also result from p orbitals overlapping end-to-end or a p orbital overlapping with any of the hybrid orbitals. Any of the hybrid orbitals can also overlap with each other end-to-end.

Any combination of two orbitals that overlap end-to-end such as these will result in a 𝜎 bond. But if we have two orbitals overlapping side-by-side, like these two p orbitals, the resulting bond is a 𝜋 bond, not a 𝜎 bond. So answer choice (C), a p orbital overlapping with another p orbital side-by-side, would result in the formation of a 𝜋 bond not a 𝜎 bond. So it’s the correct answer.

Now, let’s wrap up this video with the points that are really important. Atomic orbitals merge together to form hybrid orbitals. 𝜎 bonds are formed from the end-to-end overlap of orbitals, while 𝜋 bonds are formed by the side-by-side overlap of adjacent p orbitals. Single bonds are made of one 𝜎 bond. Double bonds are made of one 𝜎 bond and one 𝜋 bond. And triple bonds are made of one 𝜎 bond and two 𝜋 bonds.