Video transcript Voiceover: In this video, we're going to look at the SP three hybridization present in methane and ethane; let's start with methane. So that's CH four, if I want to draw a dot structure for methane, I would start with carbon, and its four valence electrons, and then we would put hydrogen around that; each hydrogen has one valence electron, so we go ahead and draw in our hydrogens with one valence electron, and that gives us the Lewis Dot structure.
Usually you see it drawn like this, with carbon with its four bonds to hydrogen around it, like that, and in methane, all of these bonds are equivalent, in terms of things like bond length and energy. And so the four valence electrons that carbon brought to the table over here, let me go ahead and highlight those four valence electrons, those should be equivalent, and if we look at the electron configuration for carbon, let's go ahead and do that right now.
It's one S two, so go ahead and put in two electrons in the one S orbital, two S two, go ahead and put in two electrons in the two S orbital, and then two P two, and so, I'm assuming you already know your electron configuration, so it would look something like that. If we look at those four valence electrons on our orbital notation here, that would be these four electrons here, the valence electrons in the outer shell.
And if we look at this, this implies that carbon would only form two bonds, because I have these unpaired electrons right here, and everything's of different energies, and so, what we see from the dot structure and experimentally, doesn't quite match up with the electron configuration here, and so to explain this difference, Linus Pauling came up with the idea of hybridization.
And so, the first thing that he said was, you could go ahead and take out one of these electrons in the two S, and promote it up to the P orbital here, so let me go ahead and show that, so we've moved one of those electrons up to the two P orbital, so we're in the excited state now.
And now we have the opportunity for carbon to form four bonds, however, those electrons are not of equivalent energy, and so Linus Pauling said, "Let's do something else here: "Let's go ahead and promote the two S orbital," so we're gonna take this S orbital, and we're gonna promote it in energy, and we're going to take these P orbitals and demote them in energy, so we're gonna lower those P orbitals, like that, so we have our P orbitals here.
And these had one electron in each of them, but we're gonna hybridize them, so this is no longer going to be an S orbital; it's going to be an SP three hybrid orbital; this is no longer going to be a P orbital; it's going to be a SP three hybrid orbital, and same with these. So the idea is, you're taking some of the S character, and some of the P character, and you're hybridizing them together into brand new orbitals, and since you're taking this from one S orbital and three P orbitals, we're doing this using one S orbital and three P orbitals, we call this SP three hybridization, so this is SP three hybridization: We create four new, hybrid orbitals.
And now we have what we're looking for, because now we have four unpaired electrons, so carbon can form four bonds now, and they're equal in energy, so that's the idea of hybridization. All right, let's think about the character, or the shape of this new hybrid orbital.
Well, we know that an S orbital is shaped like a sphere, so we're taking one of those S orbitals here, so one S orbital, and we know that a P orbital is shaped like a dumbbell, so we're taking three of these P orbitals here.
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Genetics Video Lessons. Biochemistry Video Lessons. The geometry of the sp hybrid orbitals is linear, with the lobes of the orbitals pointing in opposite directions along one axis, arbitrarily defined as the x-axis see Figure 7.
Each can bond with a 1s orbital from a hydrogen atom to form the linear BeH 2 molecule. Figure 7. The process of sp hybridization is the mixing of an s orbital with a single p orbital the pxorbital by convention , to form a set of two sp hybrids. The two lobes of the sp hybrids point opposite one another to produce a linear molecule.
Other molecules whose electron domain geometry is linear and for whom hybridization is necessary also form sp hybrid orbitals. Examples include CO 2 and C 2 H 2 , which will be discussed in further detail later. First a paired 2s electron is promoted to the empty 2p y orbital see Figure 8. This is followed by hybridization of the three occupied orbitals to form a set of three sp 2 hybrids, leaving the 2p z orbital unhybridized see Figure 9. The geometry of the sp 2 hybrid orbitals is trigonal planar, with the lobes of the orbitals pointing towards the corners of a triangle see Figure 9.
Each can bond with a 2 p orbital from a fluorine atom to form the trigonal planar BF 3 molecule. The process of sp 2 hybridization is the mixing of an s orbital with a set of two p orbitals p x and p y to form a set of three sp 2 hybrid orbitals.
Each large lobe of the hybrid orbitals points to one corner of a planar triangle. Other molecules with a trigonal planar electron domain geometry form sp 2 hybrid orbitals. Ozone O 3 is an example of a molecule whose electron domain geometry is trigonal planar, though the presence of a lone pair on the central oxygen makes the molecular geometry bent. The hybridization of the central O atom of ozone is sp 2.
Only read the boron section. Skip to main content. Covalent Bonding. Search for:. Hybrid Orbitals Learning Objectives Define hybridization. Describe sp 3 hybridization and covalent bond formation. Do you recognize this plant? Figure 1. Orbital configuration for carbon atom.
Now that carbon has four unpaired electrons it can have four equal energy bonds. The hybridization of orbitals is favored because hybridized orbitals are more directional which leads to greater overlap when forming bonds, therefore the bonds formed are stronger. This results in more stable compounds when hybridization occurs. The next section will explain the various types of hybridization and how each type helps explain the structure of certain molecules. The frontal lobes align themselves in the manner shown below.
In this structure, electron repulsion is minimized. Hybridization of an s orbital with all three p orbitals p x , p y , and p z results in four sp 3 hybrid orbitals. This Because carbon plays such a significant role in organic chemistry, we will be using it as an example here. Carbon's 2s and all three of its 2p orbitals hybridize to form four sp 3 orbitals. These orbitals then bond with four hydrogen atoms through sp 3 -s orbital overlap, creating methane.
The resulting shape is tetrahedral, since that minimizes electron repulsion. Lone Pairs: Remember to take into account lone pairs of electrons. These lone pairs cannot double bond so they are placed in their own hybrid orbital. This is why H 2 O is tetrahedral.
We can also build sp 3 d and sp 3 d 2 hybrid orbitals if we go beyond s and p subshells. The frontal lobes align themselves in the trigonal planar structure, pointing to the corners of a triangle in order to minimize electron repulsion and to improve overlap. The remaining p orbital remains unchanged and is perpendicular to the plane of the three sp 2 orbitals.
Hybridization of an s orbital with two p orbitals p x and p y results in three sp 2 hybrid orbitals that are oriented at o angle to each other Figure 3. Sp 2 hybridization results in trigonal geometry. In aluminum trihydride, one 2s orbital and two 2p orbitals hybridize to form three sp 2 orbitals that align themselves in the trigonal planar structure.
The three Al sp 2 orbitals bond with with 1s orbitals from the three hydrogens through sp 2 -s orbital overlap.
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