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Video transcript

in this series of videos we're going to look at the aromaticity or aromatic stabilization we've already seen that bromine will add a cross 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 and so benzene is an aromatic molecule let's look at the criteria to determine if a compound is aromatic alright 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 hoople'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 6 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 go ahead and sketch in the unhybridized free p orbital on each of the 6 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 6 atomic orbitals in benzene according to mo Theory those six Comic orbitals are going to cease to exist and we will get six molecular orbitals instead and 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 and your frost circle you always start at the bottom so we're going to start down here so we're going to inscribe hexagons so I'm going to see if we can put a hexagon in here all right so we have a six-sided figure 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 right so these are anti bonding molecular orbitals which are the highest and energy but 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 right so if I go back over here remember that benzene has six 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 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 so four more pi electrons and I go ahead and put those in and I have filled the bonding molecular orbitals of benzene so I have represented all six pi electrons if I think about who calls rule right 4n plus 2 I have six PI electrons so if n is equal to one who calls rule is satisfied right because I would do four times one plus two and so I would get a total of six PI electrons and so six PI electrons follows who calls rule if we look at the frost circle and we look at the molecular orbitals we can understand who calls rule a little bit better visually so if I think about if I think about these two electrons down here all right you can think about that's where the two comes from in equals rule and I think about these four electrons up here right that would be four electrons times our positive integer of one so four times one plus two 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 and so benzene is aromatic it follows our different criteria and the next few in the next videos we're going to look at several other examples of aromatic compounds and ions