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Resonance structures

Introduction to resonance structures, when they are used, and how they are drawn. Created by Jay.

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  • mr pink red style avatar for user gKaushik98
    this may seem stupid.. but,in the very first example in this video..isnt the resonating structure the same as the original?its just the inverted form of it....
    (76 votes)
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    • mr pink red style avatar for user Daniela Félix
      The equivalent ressonance structures seem like the same but there are non equivalent ressonance strutures that occur when the delocalization of electrons is between qualitativity different bonds (they are different because they bond different atoms for instance a nitrogen and a carbon and two carbons)
      (6 votes)
  • male robot hal style avatar for user apattnaik1998
    why delocalisation of electron stabilizes the ion
    (25 votes)
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    • female robot amelia style avatar for user poorvabakshi21
      Another way to think about it would be in terms of polarity of the molecule. If you have electrons that are localised on one particular atom, there would be a lot of polarity, thus the molecule would be more likely to both react and bond with other molecules. This decreases its stability.
      (9 votes)
  • aqualine ultimate style avatar for user Dexter Loh
    I still don't get why the acetate anion had to have 2 structures? aren't they both the same but just flipped in a different orientation? why does it have to be a hybrid? from what i understand, only one oxygen should be negative since a hydrogen nucleus left the molecule but what i'm seeing is that 2 oxygens are negative and this doesn't make sense
    (9 votes)
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    • piceratops ultimate style avatar for user Matt Stefely
      The two oxygens are both partially negative, this is what the resonance structures tell you!
      While both resonance structures are chemically identical, the negative charge is on a different oxygen in each. This is important because neither resonance structure actually exists, instead there is a hybrid. The oxygens share the negative charge with each other, stabilizing it, and reducing the charge on either atom.
      (26 votes)
  • marcimus pink style avatar for user eman
    Where is a free place I can go to "do lots of practice?" ... Where can I get a bunch of example problems & solutions?
    (13 votes)
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  • blobby green style avatar for user jshobikaa
    Are two resonance structures of a compound isomers??
    (5 votes)
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  • leaf green style avatar for user LAaaca
    Around I don"t understand what does the stability of whats left have to do with the leaving H+?
    (7 votes)
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    • blobby green style avatar for user jasontabletZA
      The more stable a conjugate base is the strong the acid is due to the equilibrium favoring the forward reaction a little bit more. This is relatively speaking. A non organic example are the halides, where the iodine anion is more stable than the flourine anion leading to a difference in the pKa of HF (3.2) and HI (-10). This is carried over to resonance structures, if your conjugate base has a resonance structure it's charge is delocalised and the anion is resonance stabilised, making it's corresponding acid stronger.
      (5 votes)
  • piceratops seed style avatar for user John Van Vynck
    Why at does that oxygen have a -1 formal charge? It was my understanding that oxygen's atomic number was 8, and that particular oxygen has 7 electrons. Therefore, 8 - 7 = +1, not -1. Can anyone explain where I'm wrong?
    (4 votes)
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  • leafers seedling style avatar for user RoxanaERodriguezGarcia
    How do you find the conjugate acid? They were mentioned around but it was not explained how he knew those were the conjugate bases.
    (4 votes)
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    • leafers ultimate style avatar for user hannah
      A conjugate acid/base pair are chemicals that are different by a proton or electron pair. For instance, the strong acid HCl has a conjugate base of Cl-. Remember that acids donate protons (H+) and that bases accept protons. So each conjugate pair essentially are different from each other by one proton. There's a lot of info in the acid base section too!
      (4 votes)
  • blobby green style avatar for user ian.j.alford
    I'm confused at the acetic acid briefing...

    So a single bond naturally takes only one electron from the oxygen, but then a double bond takes two more electrons? i thought it should only take one more.
    (4 votes)
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  • starky ultimate style avatar for user Vishnu Bhagyanath
    Do only multiple bonds show resonance ?
    (2 votes)
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Video transcript

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.