Chemical Bonding: Fundamental Concepts Resonance Structures and Formal Charge Electronegativity, Formal Charge and Resonance Page [1 of 3] In this lecture we’re going to pull together ideas about formal charge, resonance structures, electronegativity, and really make some predictions and some rationalizations about why molecules behave the way they do. And the first one I want to do is to go back and look at the cyanate ion. Cyanate is NCO. And in a previous lecture I talked about the fact that there were several different possible resonance structures that are all in equivalence, and that it was possible at least to figure out which one contributed the most and which one contributed the least based on formal charge.

So here are the formal charge evaluations for this left-hand resonance structure. Nitrogen has a formal charge of zero. Carbon has a formal charge of zero. Oxygen has a formal charge of -1. For B, nitrogen is -1, carbon is zero, oxygen is zero. And for C, nitrogen is -2, carbon is zero, and oxygen is +1.

Now, the rule said that formal charges of plus or minus 1 and zero are okay. In fact, zero is great. And plus or minus 2 and bigger, that’s just not going to work. Why? Because remember, formal charges reflect how the electrons are distributed relative to how they are distributed in the free atom. So how they’re distributed in a molecule versus how they’re distributed in the free atom. And in this case nitrogen has a lot more electrons than it would if were a free atom formally; in other words, an accounting method. And so this turns out not to play a particularly large role. In fact, we’ll just say it doesn’t contribute. What does it mean that it doesn’t contribute? What it means is that the distance between the carbon and the oxygen, which is predicted by this resonance structure to be very short because this is a triple bond—that’s not observed. If that’s not observed then it suggests that this picture is not valuable in our modeling of what the molecule looks like.

Okay, so now let’s take it a little bit further. A and B both contribute because we have lots of zeroes, which are really good. C doesn’t contribute at all. Both A and B contribute significantly. A contributes more because it places the negative formal charge on the more electronegative atom. What does that mean? Well, the difference between A and B is that in A, oxygen has the -1 formal charge, and in B it’s the nitrogen that has the -1 formal charge. And oxygen is more electronegative. It has a higher electronegativity so it likes the extra electrons more. And so structure A plays the biggest role. And if we had to guess, we’d say that the nitrogen-carbon bond probably looks a lot like a nitrogen-carbon triple bond rather than a nitrogen-carbon double bond. Is it likely to be a little bit longer than a conventional carbon-nitrogen triple bond? Absolutely, because of the contribution of B to the full, complete picture of what the cyanate ion looks like.

All right, let’s look at fulminate. Remember, fulminate was an isomer of cyanate, meaning that the atoms were in a different order. And if we go through the exact same exercise, the numbers here are -1, +1, -1; zero, +1, -2; and +1, +1, -3. Now, twos and threes, uh-uh. So that means B and C do not contribute. Really, A is the only resonance structure that is meaningful when we’re describing the electron distribution and any predictions we might make about the fulminate ion.

Now I remind you that mercury fulminate is the compound that I said was really unstable and was used in blasting caps. You use it to detonate dynamite, to initiate dynamite. Now the reason is that even though A is the only one that contributes, if we look at how the formal charges are distributed we notice that there’s a -1 on carbon. Carbon doesn’t have a really high electronegativity, meaning that it doesn’t really like a lot of extra charge. And yet because of the way that we put this together—and you’ll have to convince yourself or take my word for it, there’s no other way to do this—carbon always has a -1 on it. That’s not so good. Similarly, nitrogen is relatively electronegative, meaning that it likes negative charge. It’s not as electronegative as oxygen so the oxygen is really pretty happy, but the nitrogen would probably rather have a -1 than a +1 and here it has a +1. And carbon would probably rather have a +1 and here it has a -1. What does that say? It says that we can rationalize why Fulminate blows up. It’s because it’s not a particularly happy molecule. The way the electrons are distributed isn’t real good for Fulminate.

As an exercise think about the Azide anion, which is -. Sit down and go through this exact same exercise with Azide and what you’ll see is that once again we have to put a +1 on a nitrogen and that’s not so good. Nitrogen, again, is relatively electronegative. It would probably rather have a -1. It would rather have a zero for sure but if it has to have a charge, it would probably rather have a -1. Well anyway, Azide is the explosive that’s used in airbags. Sodium Azide explodes, makes , and that’s what fills the airbag when you get into a crash. So Azide is another molecule that is relatively unstable; it decomposes. And we can at least rationalize that. It doesn’t necessarily explain everything about it but we can rationalize it based on this idea that we’re putting formal charges on atoms that they don’t particularly like.

Chemical Bonding: Fundamental Concepts Resonance Structures and Formal Charge Electronegativity, Formal Charge and Resonance Page [2 of 3] Now another place where formal charge comes in is that we haven’t really talked about why we need to expand the octet, for instance on sulfur when we talk about sulfur dioxide. So here are two resonance structures for sulfur dioxide, and I can see that I don’t even need a bracket so forget about this little bracket here. And plus and minus 1, that’s not so bad, right? We said that plus and minus 1 is okay. But notice what happens when we expand the octet. When we allow 10 electrons around sulfur then what happens is the formal charges actually go from plus and minus 1 to zero, zero, and zero. And remember, zero, zero, zero, that’s the best. So we expand the octet around sulfur; what it says is that there is going to be more double bond character between the sulfur and the oxygen than might be predicted based on just these two resonance structures. And when we go to , sulfur trioxide, it’s even more important that we consider these expanded octets because for sulfur trioxide, where we just satisfy the octet rule without expanding the octet, the formal charge on sulfur has to be +2, whereas if we expand the octet it at least gets it down to plus and minus 1.

Now this exact same idea is seen in sulfate, which is eventually the conjugate base of sulfuric acid after you remove two protons. And once again, if we just do the Octet Rule we have a formal charge of +2 on sulfur, and if we expand the octet all the way up to 12 on sulfur we can actually get the formal charge of sulfur all the way down to zero. Now the minus charges are living on oxygens; that’s okay because oxygen is relatively electronegative to begin with. So this is a much better description probably of what sulfate dianion looks like, compared to this where we have the really high formal charge. And again the prediction is that there are a lot of multiple bond characters or that the sulfur- oxygen bond is going to be somewhat shorter.

Now I’ll point out that there are a lot of different resonance structures for this ion where the double bond is between the various different pairs of oxygens, so the double bond as I drew it is between this one and this one, but the double bond could be between this one and this one, or this one and this one, or this one and this one, blah, blah, blah. There are many, many different resonance structures that you could draw for the sulfate dianion. We’ll come back to that actually later on.

Now I also want to mention Exception #4. So I’ve been giving you systematically exceptions to the Octet Rule, and here’s another exception to the Octet Rule, and that is that some molecules have the best Lewis Dot Structure that doesn’t even have an octet. So let me say from the get-go, Octet Rule. Stick to Octet Rule as best you can, and then there are exceptions and I’m going to talk about those. But if you can’t remember this part, don’t worry too much because these are special cases again. Don’t let me mislead you into thinking that we don’t have to worry about the Octet Rule anymore.

If we draw boron trifluoride this way where we satisfy the Octet Rule for all the atoms, we put a positive formal charge on Fluorine and a negative formal charge on boron. And boron is a relatively electropositive element so it doesn’t really like a -1. But more importantly, Fluorine is a very electronegative element. It really likes negative charge. It really hates the fact that it has a formal positive charge. So this is not a particularly good structure. There are two more resonance structures where we put the positive charge on each of the other Fluorines but still it means putting positive formal charge on fluorine; that’s no good. But what happens if we relax the idea of having an octet? In particular, we only have to relax the idea of having an octet around boron. If we do that, maintaining an octet around fluorine, then the formal charge on boron goes to zero. The formal charge, more importantly, on fluorine goes to zero. And so what scientists will typically say—and this is actually open for some debate—is that this is a better representation of what the boron trifluoride molecule looks like; that we don’t worry about the Octet Rule, and by doing that we can make the formal charges on fluorine go to zero.

Now, inorganic chemists or people who teach general chem really make a big deal about the fact that this is an exception, but in fact in organic chemistry this idea of incomplete octets is very important in understanding reactivity in organic molecules. So here is a Lewis Dot Structure for acetone, nail polish remover. And take my word for it that the formal charge on this carbon is zero and the formal charge on this oxygen is zero. But oxygen has a large electronegativity. Remember, it really likes electrons. So suppose we redistribute the electrons in this double bond and actually put them on oxygen, and that gives oxygen a formal charge of -1 and the carbon a formal charge of +1. Remember, these formal charges are about how the electrons are really distributed in the molecule. Well, oxygen really likes that minus charge and carbon really likes the positive charge, and it turns out that this resonance structure, while not the main resonance structure—so again, Octet Rule rules, but if we consider this sort of at a higher level of theory, what it allows us to do is predict the reactivity of acetone. In other words, something that’s going to react with

Chemical Bonding: Fundamental Concepts Resonance Structures and Formal Charge Electronegativity, Formal Charge and Resonance Page [3 of 3] acetone, if it wants to react with a site of positive charge it’s going to react at the carbon, and if it wants to react with a site of excess negative charge it’s going to react with the oxygen up here.

And while we’re talking about organic molecules, let me make two final points about formal charge and resonance structure and electronegativity and that is, we can rationalize the relative acidity of alcohols versus carboxylic acids. So this is ethanol. If you’re old enough to drink, you’ve probably had an alcoholic drink at some point and unless it had lemon or vinegar or something in it, you probably wouldn’t have said that it was sour. And remember sour is characteristic of things that are acids. And so ethanol is not usually considered an acid, whereas acetic acid, which is the acid in vinegar, is clearly very sour. So this is a representation of acetic acid. And the acidity of acetic acid, which is this molecule, comes from the fact that the acetate anion is very stable relative to, say, the anion from ethanol, which is much less stable. And the reason is that for ethanol—here’s a Lewis Dot Structure—the formal charge on the oxygen is -1 and for the acetate anion the formal charge on this oxygen is -1, but we have a resonance structure. And the resonance structure moves the electrons around. It doesn’t move the nuclei around but it moves the formal charge around such that the formal charge gets moves to the oxygen up here instead of the oxygen here. Well, if we imagine that we have to take these two equally as important in order to talk about what’s going on, basically what we’re saying is the formal charge is really only -½ on this oxygen and -½ on this oxygen. So the formal charge is distributed or delocalized over a couple of sites and that turns out to be a really good situation for increasing the acidity of the conjugate acid. In other words, making this conjugate base more stable, that’s a good thing.

And this idea also works really well for mineral acids, inorganic acids like nitric acid and sulfuric acid. And to show you that let me bring up sulfuric acid again and remind you that these negative formal charges—the -1 on this oxygen and the -1 on this oxygen—remember because we have so many different resonance structures, the -1 could be here or it could be there or it could be there or it could be there. And so on average, the formal charge on each of the oxygens is actually only -½ and that helps us to rationalize the acidity of things like sulfuric acid. And you’ll have to take my word for it that for nitric acid there is a similar explanation.

So what I’ve shown you here is the interaction between electronegativity, formal charge, Lewis Dot Structures, and resonance structures, and we can use these really simple ideas and actually rationalize a great deal of chemistry, reactivity and behavior.