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Hybrid Orbitals explained - Valence Bond Theory | Orbital Hybridization sp3 sp2 sp - YouTube
Channel: Crash Chemistry Academy
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Hello and welcome to a video about
hybrid orbitals, often called valence
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bond theory. Developed in the 1930s by
the great chemist Linus Pauling as a
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model of bonding to understand the
three-dimensional placement of atoms in
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a molecule, and that is critical to our
understanding of the properties that molecules have.
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In this video we will
look at methane, ethene, ethyne, ammonia,
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and water as our models for
hybridization. There is only a small
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group of atoms in the second period that
the model really works for, but among
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those are carbon nitrogen and oxygen,
which make up the vast majority of
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molecules that exist on earth. So the
model applies to a limited number of
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elements but it applies by far to the
majority of molecules. Let's take a look
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at the word hybrid: it is a blending of
two varieties. If you get a horse and
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donkey together you get a mule, a
blending or hybrid of a horse and a donkey.
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Let's get to hybrid orbitals
using carbon as our model. As a single
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atom not bonded to anything, carbon has
two electrons in 1s, two electrons and
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2s, and two electrons in 2p. This
electron configuration is the energy
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arrangement of carbons electrons. However
carbon rarely exists in nature as an
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individual atom except momentarily while
undergoing chemical reactions.
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Carbon exists with its valence electrons bonded
to other atoms. When carbon bonds to four
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other atoms, carbon's four bonds are
experimentally seen to be equivalent. And
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so when the carbon atoms finds itself in
a bonding situation, its bonding
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electrons themselves exist at equivalent
energies, which requires that they
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hybridize to an energy that is
intermediate between the 2s energy and
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the 2p energy. Or you can think of it as
a blending or hybridization of the two
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energies, the s and p energies.
And since the energies of these
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electrons have now changed, the shape of
the orbitals they occupy are different
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as well, which we will see momentarily,
and those are called hybrid orbitals.
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They are named 2sp3 hybrid orbitals.
The naming often confuses students so
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before we go any further, let's take a
look at where the 2sp3 name comes from:
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The 2 comes from the second
principal energy level that the valence
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orbitals are in. The s comes from the 2s
orbital contributing to the
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hybridization, and the p comes from the
2p orbitals contributing to the
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hybridization, and the 3 comes from the
number of 2p orbitals used in
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hybridization. Once hybridized, the 2s and
2p orbitals no longer exist, and so we
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have four 2sp3 hybrid orbitals. Four 2sp3
hybrid orbitals derived from combining
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the energies of one 2s orbital and three 2p
orbitals, which gives a total of four 2sp3 orbitals.
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Before we look at the shape of
hybrid orbitals it would be helpful to
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briefly review the atomic orbitals. The
1s orbital is a sphere, the 2s orbital
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is a larger sphere surrounding 1s, and
here we will get rid of 1s since we
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are only concerned with the valence
electrons. Each 2s orbital is a two lobed
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shape
converging at the nucleus. So there are
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the three 2p orbitals.
However when hybridization occurs the s and
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p orbitals cease to exist, and the 2sp3
orbitals have an entirely different shape.
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We can see that orbital
hybridization explains the VSEPR
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placement of carbon's four valence
electrons since all four 2sp3
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orbitals are equivalent,
each 2sp3 orbital repels the others with
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equal force, resulting in identical bond
angles. The carbon atom only hybridizes
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when it is in a bonding situation. Here,
four hydrogen atoms bond to carbon by
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overlapping their orbitals with
carbon's hybrid orbitals. So what would be
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the reason for this? If we go back and
see that both carbon and hydrogen have
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unpaired electrons, the overlap allows
the electrons to pair and thus go to a
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lower potential energy. The illustration
here contains the valence electrons of
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both carbon and hydrogen, and since
everyone likes to visualize atoms as
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spheres we can do the same: here is our
carbon atom, and here are the hydrogens,
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with the overlapping spheres, indicating
the overlapping orbitals that constitute
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the bond. The bonds are more readily
discernible in a ball and stick model,
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which also makes the bond angle more
visible. Since all four sp3 orbitals are
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equivalent, each bonding orbital repels
the others with equal force resulting in
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identical bond angles. The bonds in
hybridization also have their own
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nomenclature. The overlapping orbitals
are called sigma bonds which represents
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the single bond occupied by a single
pair of electrons. What about double
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bonds? How does the hybridization model
explain double bonds? We will use ethene,
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C2H4, to see what happens in a double
bond. The single bonds we know are sigma
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bonds, and the double bond also has a
sigma bond, but the second bond of a
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double bond is a pi bond. Let's see how
hybridization and orbital overlap can
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explain a double bond. The two carbon
atoms in ethene are equivalent, so let's
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look at one of the carbon atoms first. A
pi bond comes from the overlap of
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unhybridized p orbitals, and so the atom
hybridizes only three orbitals,
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leaving a p orbital unhybridized for the pi bond.
The hybrid orbital is called 2sp2 the
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superscript 2 denoting that only two
2p orbitals have contributed to the
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hybridization. The 2sp2 hybrid
orbitals exist in a plane perpendicular
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to the unhybridized 2p orbital.
Let's remove the 2p orbital for now to
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more easily see that. The 2sp2 hybrid
orbitals are spread out at a 120
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degree angle, which means that
they exist in a plane, and the plane is
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perpendicular to the unhybridized 2p
orbital. So this is what both carbon
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atoms do when bonding occurs in ethene.
Each carbon atom is sp2 hybridized.
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The sigma bond occurs with 2sp2 orbital
overlap. What about the pi bond?
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The second bond of the double bond.
Previously we said that it comes from
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the unhybridized p orbitals, which we see
here from both carbon atoms. The top and
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bottom lobes of the 2p orbitals overlap
above and below the axis of the sigma
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bond forming a single pi bond. The space
in which the now paired electrons move around.
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The ball and stick model shows
this double bond with two dashes.
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To summarize, sigma bonds occur along the
axis between nuclei. The pi bond occurs
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above and below the Sigma axis where the
p orbital lobes have overlapped.
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The ethene molecule also bonds to four
hydrogen atoms by overlapping with both
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carbon's other 2sp2 hybrid orbitals,
creating four more sigma bonds. In the
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ball-and-stick model we can readily see
that each carbon has a trigonal planar
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geometry, and thus the whole molecule
exists in a plane with the single pi
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bond above and below that plane. Now
let's look at how hybridization can be a
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model for the triple bond using ethyne,
C2H2.
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The carbon-hydrogen bonds are sigma
bonds, and the triple bond is one sigma
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bond and two pi bonds. Let's see how
hybridization can accommodate this.
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Since pi bonds come from p orbitals, and we
need two pi bonds, then two 2p
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orbitals have to remain unhybridized, and
so the remaining single 2s orbital
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and a single 2p orbital will hybridize
to two 2sp hybrid orbitals. And there is
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also the violet 2px and the blue 2py
orbitals. Here each green lobe is a single
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orbital, and so they each have an
electron, while both violet lobes
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constitute the single 2px orbital with a
single electron. And both blue lobes
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constitute the single 2py orbital with
a single electron. The other carbon in
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C2H2 also has that same triple bond, and
so it has the same hybridization.
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Let's see what happens during bonding. sp
orbitals from both carbons overlap,
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forming a sigma bond. The upper lobes of
the blue 2p orbitals overlap, as do the
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lower two lobes, creating the first of
the 2 pi bonds. Can you guess where the
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second of the two pi bonds comes from?
Yes that's right!
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It is the overlapping of the violet 2p
lobes. Let's get rid of the sigma bond
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for a moment to take a look at something
interesting. Each pi bond lies on a
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separate plane and the two planes are
perpendicular to each other, and so the
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two pi bonds are perpendicular to each
other. Finally, two hydrogens will
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overlap with remaining sp hybrid
orbitals, creating the C2H2 molecule.
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The overlap of the space-filling model
reflects the overlapping orbitals, which
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is also represented by the
ball-and-stick model. In the remainder of
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the video we will look at hybridization
of nitrogen and oxygen using NH3, ammonia,
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as our model for nitrogen hybridization,
and H2O, water, for our oxygen model.
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In NH3, nitrogen has three sigma bonds and a
lone pair, so how does hybridization
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account for this? The hybridization is 2sp3, and nitrogen has 5 valence
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electrons so one of the four 2sp3 hybrid
orbitals has a pair of electrons. The three
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sigma bonds
come from the sp3 orbitals with a single
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electron, so they can pair up, and so the
remaining electron pair is a lone pair,
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an unbonded pair of electrons. As with sp3 hybridization in carbon, nitrogen
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hybrid orbitals spread out in a
tetrahedral shape. And lastly water. Here
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oxygen has two sigma bonds and two lone
pairs. In water oxygen is also 2sp3
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hybridized,
but with six valence electrons: Two of
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the sp3 orbitals have paired electrons.
You can probably guess that the sp3
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orbitals with a single electron will
overlap with hydrogen, and the remaining
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two pairs are unbonded, they are lone
pairs. Again oxygen's hybrid orbitals
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spread out in a tetrahedral shape. That's
it for hybridization, the product of a
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mad scientist! SEEYA!
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