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Studying for a test? Prepare with these 3 lessons on Chemical bonds.
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Resonance and dot structures

Video transcript
Now that we know how to draw dot structures, let's apply our rules to the nitrate anion. And we're going to see that we can draw a few different dot structures for this anion. And we're going to call those resonance structures of each other. But first, we need to calculate the total number of valence electrons. And so nitrogen is in Group 5 in the period table, therefore, five valence electrons. Oxygen is in Group 6, therefore, six valence electrons for each oxygen. I have three of them. So 6 times 3 is 18 valence electrons, plus the 5 from the nitrogen gives me 23. And I have a negative charge. This is an anion here. So we have to add one electron to that. So 23 plus 1 gives us a total of 24 valence electrons that we need to represent in our dot structure. So we know that nitrogen is going to go in the center, because oxygen is more electronegative. So nitrogen goes in the center. Nitrogen is bonded to three oxygens. So I can go ahead and put them in there like that. And let's see. How many valence electrons have we represented so far? 2, 4, and 6. Therefore, 24 minus 6 gives us 18 valence electrons left over. We're going to put those leftover valence electrons on our terminal atoms, which are our oxygens. And oxygen's going to follow the octet role. Currently, each oxygen has two valence electrons around it, the ones in magenta. So if each oxygen has two, each oxygen needs six more to complete the octet. And so I go ahead and put six more valence electrons on each one of my oxygens. Now each oxygen is surrounded by eight electrons. So the oxygens are happy. We added a total of six valence electrons to three oxygens. So 6 times 3 is 18. So we've used up all of the electrons that we need to represent. And so this dot structure, so far, it has all of our valence electrons here. Oxygen has an octet. So oxygen is happy. But nitrogen does not have an octet. If you look at the electrons in magenta, there are only six electrons around the nitrogen. And so the nitrogen wants to get to an octet. And there are a couple of different ways that we could give nitrogen an octet. For example, we could take a lone pair of electrons from this top oxygen here and move them into here to share those electrons between that top oxygen and that nitrogen. So let's go ahead and draw that resulting dot structure. So we would have our nitrogen now with a double bond to our top oxygen. Our top oxygen had three lone pairs of electrons. But now it has only two, because electrons in green moved in to form a double bond. This nitrogen is bonded to an oxygen on the bottom left and an oxygen on the bottom right here. So this is a valid dot structure. We followed our steps. And we'll go ahead and put this in brackets and put a negative charge outside of our brackets like that. So that's one possible dot structure. But we didn't have to take a lone pair of electrons from the top oxygen. We could've taken a lone pair of electrons from the oxygen on the bottom left here. So if those electrons in blue moved in here, we could have drawn another dot structure which would have been equally valid. We could have shown this oxygen on the bottom left now bonded to this nitrogen, and it used to have three lone pairs. Now it has only two. And now this top oxygen is still a single bond with three lone pairs around it. And this bottom right oxygen is still a single bond with three lone pairs around it. So this is a valid dot structure as well. So let's go ahead and put our brackets with a negative charge. And then, of course, we could have taken a lone pair of electrons from the oxygen on the bottom right. So I could have moved these in here to form a double bond. And so now, we would have our nitrogen double bonded to an oxygen on the bottom right. The oxygen on the bottom right now has only two lone pairs of electrons. The oxygen at the top, single bond with three lone pairs. And then the same situation for this oxygen on the bottom left. And so this is, once again, another possible dot structure. And so these are considered to be resonance structures of each other. And the way to represent that would be this double-headed resonance arrow here. And I think when students first see resonance structures, the name implies that, in this case, the ion is resonating back and forth between these three different possible, equally valid dot structures. And that's not quite what's going on here. Each of these dot structures is an attempt to represent the structure of the ion. But they're really not the best way of doing that. You need to think about combining these three dot structures in a resonance hybrid of each other. And so let's go ahead and draw just a simple representation of a way of thinking about a resonance hybrid. So if I combined all three of my dot structures here into one picture, I had a double bond to one oxygen in each of my three resonance structures here. And so the top oxygen had a double bond in one of them, the bottom left in the middle one, and then the bottom right in the third one. So, in reality, if we take a hybrid of all those things, we could think about the electrons being delocalized or spread out among all three of our oxygens. And so instead of giving our top nitrogen-oxygen, instead of making that a double bond, we can just show some electrons being delocalized in that area, so stronger than a single bond, but not as strong as a double bond. And we could do the same thing between this nitrogen and this oxygen. So the electrons are delocalized a little bit here. It's not a double bond. It's not a single bond. And the same idea for this nitrogen-oxygen in here. And one way we know that the ion looks more like this hybrid is because of bond length. When the ion is measured in terms of the bond length, all the nitrogen and oxygen bonds are the same length. And of course, if we thought about one of these resonance structures as being the true picture of the ion-- let's say this one, for example-- that wouldn't be the case for this ion, because this double bond here, we know that would be shorter than one of these single nitrogen-oxygen bonds. And so it's actually more of a hybrid with the electrons delocalized throughout. And that's the idea of resonance structures here.