The pi bond is the "second" bond of the double bonds between the carbon atoms, and is shown as an elongated green lobe that extends both above and below the plane of the molecule. This plane contains the six atoms and all of the sigma bonds. It is important to realize, however, that the two bonds are different: one is a sigma bond, while the other is a pi bond.
The promotion of an electron in the carbon atom occurs in the same way. As with ethene, these side-to-side overlaps are above and below the plane of the molecule. All the bond lengths and strengths in methane are roughly the same. So even though the bonds are made up of different energy orbitals they make all the same type of bonds, how can this be? Well, the way we explain it is hybridization. When these sp 3 hybrid orbitals overlap with the s orbitals of the hydrogens in methane, you get four identical bonds, which is what we see in nature.
There are two ways to form sp 2 hybrid orbitals that result in two types of bonding. Again there are two ways to form sp hybrids. The first can be formed from an element with two valence electrons in its outer shell, like lithium:.
So everything I've drawn so far is a sigma bond, so that, that. Maybe I don't want to make this picture too-- so I can just put sigma bond there, sigma bond there, sigma bond there, sigma, sigma. So far I've drawn this bond, this bond, this bond, this bond, and this bond, all of those sigma bonds. So, what happens to this last p orbital for each of these guys? Well, that's going to be kind of sticking out of the plane of the Mercedes sign, is the best way I can describe it.
And let me see if I can do that in a color that I haven't done yet. Oh, maybe this purple color. So you can imagine a pure p orbital. So a pure p orbital, I'm going to need to draw it even bigger than that, actually.
A pure p orbital, it normally wouldn't be that big relative to things, but I have to make them overlap. So it's a pure p orbital that's kind of going in, maybe you can imagine, the z-axis, that the other orbitals are kind of a Mercedes sign in the x, y plane. And now you have the z-axis going straight up and down, and those bottom two have to overlap so let me draw them bigger. So it looks like that and it looks like that. And they're going straight up and down. And notice, they are now overlapping.
So this bond right here is this bond. I could've drawn them in either way, but it's that second bond. And so what's happening now to the structure? So let me make it very clear. This right here, that is a pi bond, and this right here is also-- it's the same pi bond.
It's this guy right here. It's the second bond in the double bond. But what's happening here? Well, first of all, by itself it would be a weaker bond, but because we already have a sigma bond that's making these molecules come closer together, this pi bond will make them come even closer together. So this distance right here is closer than if we were to just have a single sigma bond there. Now, on top of that, the really interesting thing is, if we just had a sigma bond here, both of these molecules could kind of rotate around the bond axis.
They would be able to rotate around the bond axis if you just had one sigma bond there. But since we have these pi bonds that are parallel to each other and they're kind of overlapping and they're kind of locked in to that configuration, you can no longer rotate. If one of these molecules rotates, the other one's going to rotate with it because these two guys are locked together.
So what this pi bond does in the situation is it makes this carbon-carbon double bond-- it means that the double bonds are going to be rigid, that you can't have one molecule kind of flipping, swapping these two hydrogens, without the other one having to flip with it.
So you wouldn't be able to kind of swap configurations of the hydrogens relative to the other side. That's what it causes. So, hopefully, that gives you a good understanding of the difference between sigma and pi bond. And if you're curious, when you're dealing with-- just to kind of make it clear, if we were dealing with ethyne, this is an example of ethene, but ethyne looks like this. You have a triple bond. And so you have each side pointing to one hydrogen.
In this case, one of these, so the first bonds, you can imagine, so these bonds are all sigma bonds. They're actually sp hybridized. Your 2s orbital only mixes with one of the p's, so these are sp hybrid orbitals forming sigma bonds, so all of these right here.
And then both of these-- let me do this in different color. Both of these are pi bonds. And if you had to imagine it, could imagine another pi bond kind of coming out of the page and another one here coming out of the page and into the page, out and into the page, and they, too, are overlapping, and you just have one hydrogen pointing out in each direction. Each contains one electron and so is capable of forming a covalent bond. The three sp 2 hybrid orbitals lie in one plane, while the unhybridized 2p z orbital is oriented perpendicular to that plane.
The bonding in C 2 H 4 is explained as follows. One of the three sp 2 hybrids forms a bond by overlapping with the identical hybrid orbital on the other carbon atom. The remaining two hybrid orbitals form bonds by overlapping with the 1s orbital of a hydrogen atom.
Finally, the 2p z orbitals on each carbon atom form another bond by overlapping with one another sideways. It is necessary to distinguish between the two types of covalent bonds in a C 2 H 4 molecule.
A sigma bond bond is a bond formed by the overlap of orbitals in an end-to-end fashion, with the electron density concentrated between the nuclei of the bonding atoms. A pi bond bond is a bond formed by the overlap of orbitals in a side-by-side fashion with the electron density concentrated above and below the plane of the nuclei of the bonding atoms.
The figure below shows the two types of bonding in C 2 H 4. The sp 2 hybrid orbitals are purple and the p z orbital is blue. Three sigma bonds are formed from each carbon atom for a total of six sigma bonds total in the molecule. This plane contains the six atoms and all of the sigma bonds. It is important to realize, however, that the two bonds are different: one is a sigma bond, while the other is a pi bond.
Ethyne C 2 H 2 is a linear molecule with a triple bond between the two carbon atoms see Figure 4. The hybridization is therefore sp.
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