Resonance structures

Voiceover: Sometimes one dot structures is not enough to completely describe a molecule or an ion, sometimes you need two or more, and here's an example: This is the acetate anion, and this dot structure does not completely describe the acetate anion; we need to draw another resonance structure. And so, what we're gonna do, is take a lone pair of electrons from this oxygen, and move that lone pair of electrons in here, to form a double-bond between this carbon and that oxygen. And at the same time, we're gonna take these two pi electrons here, and move those pi electrons out, onto the top oxygen. So let's go ahead and draw a resonance, double-headed arrow here, and when you're drawing resonance structures, you usually put in brackets. And let's go ahead and draw the other resonance structure. So now, there would be a double-bond between this carbon and this oxygen here. This oxygen on the bottom right used to have three lone pairs of electrons around it, now it only has two, because one of those lone pairs moved in, to form that pi bond. The oxygen on the top used to have a double-bond, now it has only a single-bond to it; and it used to have two lone pairs of electrons, and now it has three lone pairs of electrons. That gives the top oxygen a negative-one formal charge, and make sure you understand formal charges, before you get into drawing resonance structures, so it's extremely important to understand that. All right, so next, let's follow those electrons, just to make sure we know what happened here. So, these electrons in magenta moved in here, to form our pi bond, like that, and the electrons over here, in blue, moved out, onto the top oxygen, so let's say those electrons in blue are are these electrons, like that. So, we have two resonance structures for the acetate anion, and neither of these structures completely describes the acetate anion; we need to draw a hybrid of these two. And so, if we take a look at, let's say the oxygen on the bottom-right here, we can see there's a single-bond between this carbon and this oxygen. If we look at this one over here, we see there is now a double-bond between that carbon and the oxygen. So, if you think about a hybrid of these two resonance structures, let's go ahead and draw it in here, we can't just draw a single-bond between the carbon and that oxygen; there's some partial, double-bond character there. So, it's a hybrid of the two structures above, so let's go ahead and draw in a partial bond here, like that. The exact same thing for the top oxygen: Here we have a double-bond, and then over here we have a single-bond, so somewhere in between is going to be our hybrid. So let's go ahead and draw that in. So, we can't just draw a single-bond in our hybrid; we have to show some partial, double-bond character, drawing the dotted line in there, like that. And also charge, so if we think about charge, the negative charge is on the oxygen on the bottom-right, and then over here the negative charge is on the top oxygen. And, so that negative charge is actually de-localized, so it's not localized to one oxygen; it's de-localized, it's distributed evenly, over both of those oxygens, here. And so this is just one way to represent the hybrid, here, and studies have shown that the hybrid is closer to what the actual anion looks like. So, studies have been done on these bond lengths here, and the bond between this carbon and this oxygen, it turns out to be the exact same bond length as the bond between the carbon and this oxygen, so, it's the exact same bond length. And so, the hybrid, again, is a better picture of what the anion actually looks like. Let's think about what would happen if we just moved the electrons in magenta in. So if I go back to the very first thing I talked about, and you're like, "Well, why didn't "we just stop, after moving these electrons in magenta?" Let's go ahead and draw what we would have, if we stopped after moving in the electrons in magenta. So we would have this, so the electrons in magenta moved in here, to form our double-bond, and if we don't push off those electrons in blue, this might be our resonance structure; the problem with this one, is, of course the fact that this carbon here has five bonds to it: So, one, two, three, four, five; so five bonds, so 10 electrons around it. We know that carbon can't exceed the octet of electrons, because of its position on the periodic table, so this is not a valid structure, and so, this is one of the patterns that we're gonna be talking about in the next video. So the pattern is, a lone pair of electrons, so next to a pi bond, which is the example we see here for the acetate anion, and so these are the two resonance structures. The problem with the word, "resonance," is, when you're a student, you might think that the anion will resonate back and forth between this one and this one; that's just kind of what the name seems to imply. And that's not actually what's happening; it's just that we can't draw, if we're just drawing one dot structure, this is not an accurate description, and so the electrons are actually de-localized, so it's not resonating back and forth. When you draw resonance structures in your head, think about what that means for the hybrid, and how the resonance structures would contribute to the overall hybrid. So don't forget about your brackets, and your double-headed arrows, and also your formal charges, so you have to put those in, when you're drawing your resonance structures. All right, let's look at an application of the acetate anion here, and the resonance structures that we can draw. If we look at the acetate anion, so we just talked about the fact that one of these lone pairs here, so this is not localized to the oxygen; it's de-localized, so we can move those electrons in here, we push those electrons off, onto the oxygen, we can draw a resonance structure, and so this negative-one formal charge is not localized to this oxygen; it's de-localized. And so, because we can spread out some of that negative charge, that increases the stability of the anion here, so this is relatively stable, so increased stability, due to de-localization. So, the fact that we can draw an extra resonance structure, means that the anion has been stabilized. And so, this is called, "pushing electrons," so we're moving electrons around, and it's extremely important to feel comfortable with moving electrons around, and being able to follow them. So, the only way to get good at this is to do a lot of practice problems, so please do that; do lots of practice problems in your textbook. If we compare that to the ethoxide anion, so over here, if we try to do the same thing, if we try to take a lone pair of electrons on this oxygen, and move it into here, we can't do that, because this carbon right here, already has four bonds; so it's already bonded to two hydrogens, and then we have this bond, and this bond. And so, moving those electrons in, trying to de-localize those electrons, would give us five bonds to carbon, and so we can't do that; we can't draw a resonance structure for the ethoxide anion. So those electrons are localized to this oxygen, and so this oxygen has a full, negative-one formal charge, and since we can't spread out that negative charge, or it's going to destabilize this anion. So this is not as stable, so decreased stability, compared to the anion on the left, because we can't draw a resonance structure. If we think about the conjugate acids to these bases, so the conjugate acid to the acetate anion would be, of course, acetic acid. So we go ahead, and draw in acetic acid, like that. The conjugate acid to the ethoxide anion would, of course, be ethanol. So we go ahead, and draw in ethanol. And we think about which one of those is more acidic. We know that acetic acid is more acidic; it's more likely to donate a proton, because the conjugate base is more stable, because, you could think about resonance, or de-localization of electrons. If you're looking at ethanol, ethanol's not as likely to donate its proton, because the conjugate base, the ethoxide anion is not as stable, because you can't draw any resonance structures for it. The negative charge is not able to be de-localized; it's localized to that oxygen. So this is just one application of thinking about resonance structures, and, again, do lots of practice. In the next video, we'll talk about different patterns that you can look for, and we talked about one in this video: We took a lone pair of electrons, so right here in green, and we noticed this lone pair of electrons was next to a pi bond, and so we were able to draw another resonance structure for it. We don't have that situation with ethoxide: We have a lone pair of electrons, but we don't have a pi bond next to it, And so, more in the next video on that.