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Resonance

Resonance arises when more than one valid Lewis structure can be drawn for a molecule or ion. The overall electronic structure of the molecule or ion is given by the weighted average of these resonance structures and is referred to as the resonance hybrid. Created by Sal Khan.

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  • starky seed style avatar for user Runtian Du
    Does covalent bond have strict orientations?
    Is all the different resonance just different rotations of the same molecule since if you take one and rotate 60 degrees you would get the other ones.
    (13 votes)
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    • male robot donald style avatar for user achawla
      Think about it this way, the 'other ones' that you mentioned are just alternate ways of drawing the lewis structure of the covalent bond, both of which are equally likely, thus, in reality the molecule actually exists in equilibrium with all three with partial bonds. For example, NO3- resonance structures that sal shows are all equally likely because like you said you could rotate them and get the other ones, or in more scientific terms, have equally favorable formal charges and in their are actually 1 1/3 bonds between the nitrogen and each of the three oxygens.
      (1 vote)
  • leafers ultimate style avatar for user Nick Voelckers
    At until , Sal draws the resonance hybrid Lewis structure for the Nitrate anion. Contrary to his other Lewis structures for nitrate, he didn't add the valence electrons around the oxygen atoms. When drawing these kinds of structures, are the valence electrons included? If so, how are they drawn (do the dotted line covalent bonds take up a lone pair for each oxygen)?
    (9 votes)
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  • piceratops ultimate style avatar for user MEHER
    @ what does sal actually mean by saying regards to different ways of visualizing the nitrate anion, "these contribute to a resonance hybrid" what do u mean by 'contribute' here in this context??
    (2 votes)
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    • leaf red style avatar for user Richard
      Certain molecules cannot be represented by a single Lewis structure. Instead there are multiple possible, valid structures which could exist which we call resonance structures. The question then becomes, which one of these resonance structures represents the true structure of the molecule? The answer is all of them do. The true structure is a combination of all the resonance structures which we refer to as the resonance hybrid.

      A second question arises, do the constituent resonance structures contribute equally to the resonance hybrid? The answer is, no, certain resonance hybrid contribute to the molecule’s structure to a greater degree than others. The more stable resonance structures contribute more so to the resonance hybrid than do the less stable ones. Stable resonance structures features include having fulfilled octets and absent formal charges. Or if we have to have formal charge, placing the negative ones on more electronegative atoms and positive ones on less electronegative atoms.

      Hope that helps.
      (7 votes)
  • blobby green style avatar for user krishanthsenthil
    I dont understand. oxygen has 6 electrons how did sal add 6 electrons and a bond pair to each oxygen?
    (3 votes)
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  • blobby green style avatar for user Manahil Kashif
    How can you be in between a single bond and a double bond? is it like only sharing one electron? And what does it mean for electrons to be delocalized throughout?
    (3 votes)
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  • blobby green style avatar for user gupick gupick
    why does nitrogen have 4 bonds if it only has 5 valence electrons?
    (2 votes)
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    • leaf red style avatar for user Richard
      The valance electrons of a single atom are simply the number of electrons that atom brings to the molecule. Once all the bonds form, there’s no longer a sharp distinction between whose electrons are whose. All the valance electrons are essentially collectively managed. Both the nitrogen and the oxygen are following the octet rule which means they are trying to form bonds so they can feel like they have eight valence electrons. So nitrogen having four bonds is its way of following the octet rule (each bond has two electrons so 4x2 = 8).

      Hope that helps.
      (2 votes)
  • male robot hal style avatar for user PUDIFIRE
    so in actuality how does a resonance hybrid exist like rather than an electron staying between 2 atoms in a molecule is it bouncing off in the entire molecule?
    (1 vote)
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    • leaf red style avatar for user Richard
      In a resonance hybrid we have delocalized electrons. Delocalized electrons are electrons which are not solely bound to a single atom or bond, but rather exist spread out over several atoms in the molecule. As opposed to localized electrons which do remain bound to a single location in the molecule.

      So the resonance hybrid is a single structure averaged from the individual resonance structures where the delocalized electrons spread out across the molecule and often help to distribute charge to several atoms. The also results in bonds which have noninteger values. For example, in nitrate in the video, the nitrogen-oxygen bonds are not single or double bonds, but rather 1.3 bonds after averaging the three resonance structures.

      Hope that helps.
      (3 votes)
  • marcimus pink style avatar for user 25189
    how do you know its resonance
    (2 votes)
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  • piceratops ultimate style avatar for user Tanzz
    At , Sal allocates two valence electrons of Oxygen to Nitrogen to make another covalent bond. Aren't covalent bonds mean the bonds with mutual sharing of electrons from all the elements? Here, the Nitrogen isn't mutually sharing its one electron with Oxygen to make a double bond.
    (1 vote)
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    • leaf red style avatar for user Richard
      Well Sal only put those lone pairs on each of the oxygen atoms in the first place to use up all the valence electrons according to the method used to create Lewis structures. If we think of each atom's electron belonging to it then even before we form the double there's some unequal distribution in a sense. Each oxygen brings six valence electrons as individual atoms and the nitrogen atom brings five electrons. And before the double bond is formed each oxygen has 7 electrons and the nitrogen has only 3 electrons which means we've allocated some of the nitrogen's electrons to the oxygens already.

      Basically it's more correct to think about these valence electrons being more fluid and being able to move around the molecule more easily than being static and belonging to a single atom. So when we take that lone pair from oxygen, it's not really us taking oxygen's electrons since they never completely belonged to it in the first place. And remember the most important point is that it's not us creating these bonds, it's us trying to determine what type of bonding is actually occurring the molecule.

      So when we correctly identify the bond orders (single and double bonds) then these bonds are indeed covalent bonds as you stated where the electrons are being shared between the bonding atoms.

      Hope that helps.
      (3 votes)
  • ohnoes default style avatar for user MarieLydia13
    I cannot figure out at , what do these three and two blue points next to the single bond mean?
    (2 votes)
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Video transcript

- [Instructor] Let's see if we can draw the Lewis diagram for a nitrate anion. So a nitrate anion has one nitrogen and three oxygens, and it has a negative charge. I'll do that in another color. It has a negative charge. So pause this video and see if you can draw that, the Lewis structure for a nitrate anion. All right, well we've done this many times. The first step is to just account for the valence electrons. Nitrogen has one, two, three, four, five valence electrons in its outer shell, and in that second shell, if it's a neutral, free nitrogen atom. So we have five valence electrons there. Oxygen has one, two, three, four, five, six valence electrons. But if you have three oxygens, you're going to have six times three. And so if you just add up the valence electrons, if these were free, neutral atoms, you would get five plus 18, which is 23 valence electrons. Now the next thing we have to keep in mind is this is an anion. This has a negative one charge right over here. So it's going to have one more extra electron, one more extra valence electron than you would expect if these were just free atoms that were neutral. So let's add one valance electron here. So that gets us to 24 valence electrons. And then the next step is let's try to actually draw this structure. And the way we do it is we try to pick the least electronegative atom that is not hydrogen to be the central atom. And in this case it is nitrogen. It's to the left of oxygen in that second period. So let's put nitrogen in the center, right over there. And around it let's put three oxygens. So one, two, three oxygens. Let's put a single bond between them. And so so far we, and let me do that in another color, so we can account for it better. So I'll do them in purple. So so far we have accounted for two, four, six valence electrons. So minus six valence electrons gets us to 18 valence electrons. The next step is we would try to allocate as many of these as possible to our terminal atoms, the oxygens over here. Try to get them to a full octet. So let's do that. This, each of these oxygens, they're already participating in one of these covalent bonds, so they already have two valence electrons hanging out. So let's see if we can give them each another six, to get to eight. So two, four, six. Two, four, six. And then two, four, and six. And so just like that we have allocated 18 valence electrons, six, 12, 18. So minus 18 valence electrons. And we are now left with no further valence electrons to allocate. But let's see how our atoms are doing. We know the oxygens have a full octet, but the nitrogen only has two, four, six valance electrons hanging around. It would be great if there was a Lewis structure where we could have eight valence electrons for that nitrogen. Well one way to do that is to take one of these lone pairs from one of the oxygens and turn that into another covalent bond. So let's do that. So let me just erase this pair right over here, and I'm just going to turn that into another covalent bond. And this is looking pretty good. We have eight valence electrons around each of the oxygens. And now we have eight for the nitrogen, two, four, six, eight. And we have to remind ourselves that this is an anion. It has a negative one charge. So to finish the Lewis diagram we would just put that negative charge there. And this is all well and good, but if this was the only way that nitrate existed when we observed nitrate anions in the world, we would expect to see one shorter bond and two longer bonds, and we would expect one of the bonds to have a different energy than the other two. But in the real world we don't see that. We see that all of the bonds actually have the same length, and they actually have the same energy. And so an interesting question is why is that? And one thing that you might appreciate is, when I took that lone pair to create this covalent bond, I could have done it with that top oxygen. I could have done it with this bottom-left oxygen. Or I could have done it with that bottom-right oxygen. And so there's actually three valid Lewis structures that we could have had. Not only could we have had this Lewis structure, we could have had this one, and I'll draw it all in yellow to save us some time, where you have this nitrogen. It has a single bond with that top oxygen. And so that top oxygen still has six electrons in lone pairs. And maybe it forms a double bond with the bottom-left oxygen. So this bottom-left oxygen only has two lone pairs. One of them would have gone to form the double bond. And then this oxygen would look the same. So what I am drawing here is another valid Lewis structure. Or the double bond might have formed with this bottom-right oxygen, so let me draw that. So another valid Lewis structure could look like this. So nitrogen bonded to that oxygen has three lone pairs. This oxygen also has three lone pairs. And now this one has the double bond and only has two lone pairs. And whenever we see a situation where we have three valid Lewis structures, we call this resonance. Resonance. Resonance. And we'll put an arrow, these two-way arrows between these structures. And when you hear the word resonance, it sometimes conjures up this image that you're bouncing back, you're resonating between these structures. But that's actually not right. What the right way to think about it is, these different ways of visualizing the nitrate, these contribute to a resonance hybrid, which is really the true way that the nitrate exists. And so, if we wanted to draw a resonance hybrid, it would look like this. You have the nitrogen in the center. You have your oxygens, one, two, three. I can draw our first covalent bond like that. And then you would show the bond between nitrogen and each of these oxygens are a hybrid between someplace between a single bond and a double bond. And so instead of just one of them having the double bond and the other two having single bonds, they're all somewhere in between. So maybe you draw a dotted line, something like that, to show what the reality is, is that you actually have three bonds that are someplace in between a single and a double bond, because the electrons in this molecule are delocalized throughout. And of course you wanna make sure, you always wanna make sure that people recognize that this is a anion. So this is the idea of resonance. You have multiple valid Lewis structures. They all contribute to a resonance hybrid, which is actually what we observe. We're not just bouncing between these different structures. The actual observation will be a hybrid of the three. Now what we just drew here, these three are all equivalent. But in certain cases, we'll see this in future videos, you don't have equivalent structures, and some of them might contribute more to the resonance hybrid than others. But we'll see that in future videos.