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Aromatic stability I

The aromaticity of benzene. Created by Jay.

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  • duskpin ultimate style avatar for user Hideki Minamizaki
    At , what is Atomic Orbitals and Molecular Orbitals? What is the difference between them? Is there a video covering this?
    Thank you ^-^
    (8 votes)
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  • blobby green style avatar for user Sean Zumel
    How did you know to assign 1 for N?
    (8 votes)
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  • blobby green style avatar for user Andrew Yevugah
    At , why is n=1 and not 6?
    (2 votes)
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  • leaf green style avatar for user Ana
    What does the number n in Huckel's rule represent?
    (3 votes)
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    • leafers seed style avatar for user Ikenna Frank
      n is any positive integer (whole number).
      So for example, we know for anti- aromatic compounds, 4n= number of p orbitals. The number of p orbitals in an [10] annulene (cyclodecapentane) is 10 and so we can find out that n is not an integer if we use the equation mentioned above and therefore conclude that cycldecapentane is aromatic. Assuming that the annulene is planar.
      (Annulene: hydrocarbons with alternating single and double bonds)
      (1 vote)
  • blobby green style avatar for user Julian Christopher Lee
    How do you know when a compound is aromatic, non aromatic or antiaromatic? If you follow the 2 rules it will be aromatic and if you have 4n rather than 4n +2 it will be antiaromatic but how do you know when the conformation changes to become more stable and the P orbitals are not in phase anymore making it nonaromatic?
    (2 votes)
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    • orange juice squid orange style avatar for user awemond
      This really depends on a case-by-case basis. An antiaromatic compound will change its conformation to be non-aromatic (by moving its p orbitals out of the plane) anytime the energy gained in moving to this new conformation is smaller than the energy lost by no longer being antiaromatic (ie it happens whenever there is a net energy loss and thus the changed-conformation compound is more stable). Usually, unless a compound is very conformationally restrained (for example a very small ring), you can expect there to be at least some conformational distortion to get the p orbitals out of plane if the compound is antiaromatic.
      (3 votes)
  • spunky sam blue style avatar for user Chunmun
    At , what that circle is called ? And why it is used here ?
    (1 vote)
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  • mr pants teal style avatar for user RVC.LOBO3
    how will we come to know that the given molecule is planar?
    thank u 4 answering.
    (2 votes)
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  • blobby green style avatar for user packshack
    What is the logic behind Huckel's rule and the frost circle?
    (2 votes)
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    • spunky sam blue style avatar for user Ernest Zinck
      Huckel's Rule is simply a summary of many experimental observations: a planar, cyclic, conjugated system is aromatic if it contains 4n + 2 π electrons. Similarly, In 1953, A. A. Frost noted that these systems were inscribed within a circle with a point at the bottom, the heights of the vertices on the circles reflected the relative energies of their molecular orbitals. It was all a matter of recognizing patterns in the observations.
      (1 vote)
  • leaf grey style avatar for user Steven Meaney @Atlas
    Why did Hofmann start describing in terms of smell if not all aromatics have a distinctive smell?
    (1 vote)
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    • orange juice squid orange style avatar for user Samantha Maxey
      Originally it was thought that all aromatics had that smell, however upon trying to define it other molecules were added. As the unusual stability of aromatic compounds was investigated, the term aromatic came to be applied to compounds with this stability, regardless of their odor.
      (2 votes)
  • blobby green style avatar for user Amir Luchtenstein
    i didn't understand why he drew the molecular orbitals of benzen as "bonding" and not as "anti"?
    (1 vote)
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    • blobby green style avatar for user Andrew Vermiglio
      Imagine his drawing on a graph. Energy is on a y-axis (meaning the top, anti-bonding orbital is the least stable because it has the highest potential energy).
      Electrons in molecular orbitals tend to fill the low-energy spots first. Notice the order in which he drew the electrons. That is the order you must use.
      (2 votes)

Video transcript

In this series of videos, we're going to look at aromaticity or aromatic stabilization. We've already seen that bromine will add across a double bond of a simple alkene like cyclohexene to give us a mixture of enantiomers for our products. If we try the same reaction with benzene, we're not going to get anything for our product. So there's no reaction. And so benzene is more stable than cyclohexene. At first, you might think that the stability is due to the fact that benzene is conjugated. But numerous other experiments have shown that it is even more stable than we would expect. And that extra stability is called aromaticity or aromatic stabilization. So benzene is an aromatic molecule. Let's look at the criteria to determine if a compound is aromatic. So a compound is aromatic If it contains a ring of continuously overlapping p orbitals. And so if the molecule is planar, that's what allows the p orbitals to overlap. It also has to have 4n plus 2 pi electrons in the ring, where n is equal to 0, 1 2, or any other positive integer. And this is called Huckel's rule. So let's go ahead and analyze benzene in a little bit more detail. So if I look at the dot structure, I can see that benzene has 2 pi electrons there, two here, and two more here, for a total of six pi electrons. If I look at the carbons of benzene, I can see that each carbon has a double bond to it. So each carbon is sp2 hybridized. And if each carbon is sp2 hybridized, that means that each carbon has a free p orbital. So I'm going to go ahead and sketch in the unhybridized free p orbital on each of the six carbons of benzene. Now, since benzene is a planar molecule, that's going to allow those p orbitals to overlap side by side. So you get some overlap side by side of those p orbitals. And so benzene contains a ring of continuously overlapping p orbitals. So p orbitals are considered to be atomic orbitals. And so there are a total of six atomic orbitals in benzene. According to MO theory, those six atomic orbitals are going to cease to exist. And we will get six molecular orbitals instead. So benzene has six molecular orbitals. Drawing out these molecular orbitals would be a little bit too complicated for this video. So check out your textbook for some nice diagrams of the six molecular orbitals of benzene. However, it is important to understand those six molecular orbitals in terms of their relative energy levels. And the simplest way to do that is to draw a frost circle. And so here I have a circle already drawn. And inside the circle we're going to inscribe a polygon. And since benzene is a six-membered ring, we're going to inscribe a hexagon in our frost circle. I'm going to go ahead and draw a center line through the circle, just to help out with the drawing here. And when you're inscribing your polygon in your frost circle, you always start at the bottom. So we're going to start down here. So we're going to inscribe a hexagon. Let's see if we can put a hexagon in here. And so we have a six-sided figure here in our frost circle. The key point about a frost circle is everywhere your polygon intersects with your circle, that represents the energy level of a molecular orbital. And so this intersection right here, this intersection here, and then all the way around. And so we have our six molecular orbitals. And we have the relative energy levels of those six molecular orbitals. So let me go ahead and draw them over here. So we have three molecular orbitals which are above the center line. And those are higher in energy. And we know that those are called antibonding molecular orbitals. So these are antibonding molecular orbitals, which are the highest in energy. If we look down here, there are three molecular orbitals which are below the center line. And those are our bonding molecular orbitals. So those are lower in energy. And if we had some molecular orbitals that were on the center line, those would be non-bonding molecular orbitals. We're going to go ahead and fill our molecular orbitals with our pi electrons. So go back over here. And remember that benzene has 6 pi electrons. And so filling molecular orbitals is analogous to electron configurations. You're going to fill the lowest molecular orbital first. And each orbital can hold two electrons, like electron configurations. And so we're going to go ahead and put two electrons into the lowest bonding molecular orbital. So I have four more pi electrons to worry about. And I go ahead and put those in. And I have filled the bonding molecular orbitals of benzene. So I have represented all 6 pi electrons. If I think about Huckel's rule, 4n plus 2, I have 6 pi electrons. So if n is equal to 1, Huckel's rule is satisfied. Because I would do 4 times 1, plus 2. And so I would get a total of 6 pi electrons. And so 6 pi electrons follows Huckel's rule. If we look at the frost circle and we look at the molecular orbitals, we can understand Huckel's rule a little bit better visually. So if I think about these two electrons down here, you could think about that's where the two comes from in Huckel's rule. If think about these four electrons up here, that would be four electrons times our positive integer of 1. So 4 times 1, plus 2 gives us six pi electrons. And we have filled the bonding molecular orbitals of benzene, which confers the extra stability that we call aromaticity or aromatic stabilization. And so benzene is aromatic. It follows our different criteria. In the next few videos, we're going to look at several other examples of aromatic compounds and ions.