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
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
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
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
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.