[MUSIC] In Lewis dot structures the electrons are shown to be localized. These six electrons in this triple bond are all shown to be localized between those two carbons. But some molecules can be drawn in multiple ways. Where we have different number of electrons in different places depending on how we draw it. Let me show you an example. So an example of this is a deprotonated formic acid, also known as the formate anion. [SOUND] Formate anion has one hydrogen, one carbon, and two oxygens. But it has four bonds. So I have a choice, about where I draw the double bond. Doesn't make a, very much sense to put a double bond to hydrogen, so I'm not going to do that. Hydrogen's very, very small, and it really can only accommodate a single bond. But I could put a double bond to the blue oxygen, the one that's one the lower right, or I could draw it a different way and put the double bond to the yellow oxygen, the one that's going up. The other oxygen then has an extra lone part, pair and ends up with a formal charge of negative one. But these are equivalent energy resonance structures. Both of them have one carbon with no formal charge, one hydrogen with no formal charge, one oxygen with no formal charge, and one oxygen with a negative one formal charge. So in terms of the electron density around the atoms, these two are equivalent. When we can draw a molecule in two different ways, we say we're drawing resonance structures. So here again are those two resonance structures of the formate anion, and to show that they're resonance structures, I use this double-headed arrow between them. Sometimes people will even put them into brackets. But I often leave the brackets off and so do a lot of other people, but this double-headed arrow is really important. It shows resonance. I can convert from one to the other simply by moving electrons. So resonance structures show the delocalized nature of some of the electrons. So, I can push the electrons from the yellow oxygen onto the blue oxygen. What do I mean by say push the electrons? Well I can move electrons with arrows. A double-headed arrow shows two electrons moving, [SOUND] and a single-headed arrow, which is sometimes called a fishhook, shows one electron moving. [SOUND] In this case I'm going to move two electrons at a time. I'm going to start with the electrons on the oxygen that has the formal charge, the yellow oxygen here on the left. And I'm going to move both of those electrons of the lone pair into the region between the oxygen and the carbon. So the lone pair's becoming a bonding pair. When it does that, that's going to put five bonds to the carbon. Do you see that? Carbon will not allow that. So the only way the oxygen can move those electrons down onto the oxy, on to the carbon, excuse me, is if that carbon can kick some other electrons out to make room for the incoming electrons. And that's what it does. It takes these pie bonding electrons that are to the other oxygen, and it moves those up onto the oxygen and makes them into a lone pair. Do you see that those two green arrows that I've drawn on on that structure convert the structure on the left into the structure on the right? It now has a double bond that came from where the lone va, pair was before, and it has a lone pair that used to be the double bond. So that's how we can push electrons to interconvert between two equivalent energy resonance structures. Another way we can show resonance structures is to show the dash line notation. So for formate the dash line notation would have a dash between each carbon-oxygen pair, and then each oxygen actually has a negative one-half formal charge. This is actually a little bit more realistic representation of what the molecule really looks like in real life. If you look up at the top structure, it sort of looks like the two oxygens are playing catch with the electrons, right? Throwing them back and forth. And it also looks like the bonds are oscillating between being longer and shorter, doesn't it? Because we just learned on a previous slide that a single bond is longer than a double bond, and I've tried to show that here in my picture. Do you see that this yellow bond here looks slightly longer than this double bond? And so what would be happening on the top picture, is that the bonds would be getting longer and shorter as the a, [LAUGH] as the electrons are being moved around. That little motion didn't really show it very well, but hopefully you get the idea. But in the bottom picture it's showing us more realistically that the bond order between carbon and oxygen is not one a and it is not two, we have three bonds and only two regions, so to calculate the boundary, you take the number of bonds and divide it by the number of regions. Region is just an area between two atoms, and we see that the real bond order is one and a half. So, this bond between carbon and oxygen is slightly shorter than a single bond, and slightly longer than a double bond. Each oxygen also has an equivalent charge, and it has that charge all the time. So, it's not like the oxygen's 0 for a second, and then minus 1 for second, and then back to 0. Both oxygens are minus one-half formal charge all the time. I like the dashed-line notation because it shows more accurately that the bonds are equal length. But you're going to find that in organic chemistry, we frequently show the resonance structures more like they're shown at the top. So you just need to get used to that. So what do I mean when I keep saying resonance? Resonance is a simply a way to show you the delocalization of electrons in a molecule. And it's really important to remember that the atoms are not throwing the electrons back and forth as is often shown in those Lewis dot structure interconversions. The true electron density of a molecule is actually a hybrid between the different resonance structures that can be written down on the paper. Let's learn some simple rules for resonance structure interconversion. How can we move the electrons to get from one resonance structure to another? Well the first thing you have to remember is that you can move electrons but you can't move atoms. The total number of electrons doesn't change. The total number of electrons is determined by the number of valence electrons that each atom brought to the molecule when the bonds were made. The second period elements never expand their octet. So carbon never has more than four bonds. Nitrogen and oxygen never have more than four bonds. If you get down into the third period, where sulfur and phosphorous are, then you can draw some structures that have maybe five or six bonds to the sulfur or the phosphorous, but don't do it for the second period elements. The most stable structure is the one that has minimal charge separation. So if you can draw more than one Lewis dot structure, but one of them has formal charges and the other one doesn't, the more stable structure is the one that has no formal charge. Sometimes it's easy, easiest to figure that out by actually drawing the structure a couple of different ways and then picking the best one. Sometimes it's difficult to predict which one would be the best arrangement of bonds and lone pairs. Let's do that for a molecule, let's just do an example, [SOUND] for this inorganic molecule. Now the first item that's written in this structure is usually the one you put in the middle. [SOUND] And I've calculated with some math that this molecule needs to have three bonds and six lone pairs. I have to draw in three bonds. So I can draw one bond to the oxygen and two bonds to the chlorine or I can do it the other way around. Right? I could draw two bonds to my oxygen and only one bond to the chlorine. The next thing I need to do is add in lone pairs to give everything an octet. Well, the oxygen on the left only has one bond to it, so it needs to have three lone pairs. There we go. The nitrogen in the middle has, is also short of an octet. It needs one more lone pair to make an octet. So I'll go ahead and add that. And the chlorine on the end needs two more lone pairs in order to give it an octet. If added up how many valence electrons these atoms brought from the periodic table, nitrogen had five, oxygen had six, and chlorine had seven, so that's 18 electrons. Do I have 18 electrons in my picture? Let's count. 2, 4, 6, 8, 10, 12, 14, 16, 18. [SOUND] I do have 18 electrons. So that makes me feel pretty good that, that structure isn't too bad, has the right number of electrons. Let's do the same thing for the structure on the right. Oxygen needs two more lone pairs in order for it to meet the octet rule. Nitrogen in the middle needs one more lone pair. And the chlorine over on the right now needs three lone pairs. So I'm just drawing those dots in. And once again I have 18 electrons if I count. [SOUND] So that checks out with the number of electrons I should have. But I'm not quite finished yet, am I? I need to check to see if there's any formal charge on these atoms. On the periodic table you can see that oxygen is in group six. So, if I go over to the oxygen on the left, it should have six electrons. But this oxygen formally has, let me just count with red. 1, 2, 3, 4, 5, 6, 7. [SOUND] So this oxygen has a minus 1 formal charge, doesn't it? What about nitrogen? Nitrogen should have five from the periodic table. And this nitrogen does have 1, 2, 3, 4, 5. So, 5 minus 5 equals 0. No formal charge on the nitrogen. What about the chlorine on the right? Chlorine is in group seven. So, chlorine has seven valence electrons that it's bringing with it to make the molecule. That's just from looking on the periodic table. This chlorine formally owns 1, 2, 3, 4, 5, 6. So I'm going to say 7 minus 6, which is plus 1. So this chlorine has a plus 1 formal charge. So there are some formal charges on the molecule on the left. What about the molecule on the right? Oxygen should have six electrons and this oxygen has 1, 2, 3, 4, 5, 6. Uh-huh, no formal charge. What about the nitrogen? Nitrogen should have five valence electrons. Again, I'm just looking this up on the periodic table. This nitrogen has 1, 2, 3, 4, 5. So no formal charge on the nitrogen. And chlorine should have seven valence electrons. This chlorine has 1, 2, 3, 4, 5, 6, 7. Uh-huh. The structure on the right has no formal charge. So it is the most stable structure, because it doesn't have charge separation. Remember opposite charges attract, so this positively charged chlorine over here, is very attractive to those electrons, so they're going to move towards the chlorine. How do they do that? They do that by coming down and making a bond with the nitrogen. And then the nitrogen which can't accommodate the five bonds or lone pairs around it kicks out these electrons onto the chlorine and that gives the structure on the right. So the preferred lower energy structure here, is the structure on the right. But these are two resonance structures of the same compound. It just happens that the one on the left is higher energy. Doesn't mean that it doesn't exist? And it's not illegal or anything like that. It follows all the rules. It just exists at higher energy. It's less stable. [SOUND] Which means it's more likely to react, and we're less likely to encounter it just floating around in nature. Systems in nature seek the lowest energy spontaneously, don't they? Let's look at more resonance structure interconversion rules. Some finer points. Because in organic chemistry, you'll find there's a lot of patterns that repeat over and over, and it's good to know those patterns so you can recognize when they're happening. But the devil is in the details, that's what gets you. You need to know the pattern so well that you're then able to look at the intricate differences between things.