CHEM 125a: Freshman Organic Chemistry I

Lecture 2

 - Force Laws, Lewis Structures and Resonance


Professor McBride begins by following Newton’s admonition to search for the force law that describes chemical bonding. Neither direct (Hooke’s Law) nor inverse (Coulomb, Gravity) dependence on distance will do - a composite like the Morse potential is needed. G. N. Lewis devised a “cubic-octet” theory based on the newly discovered electron, and developed it into a shared pair model to explain bonding. After discussing Lewis-dot notation and formal charge, Professor McBride shows that in some “single-minimum” cases the Lewis formalism is inadequate and salvaging it required introducing the confusing concept of “resonance.”

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Freshman Organic Chemistry I

CHEM 125a - Lecture 2 - Force Laws, Lewis Structures and Resonance

Chapter 1. Newton’s “Additions”: An Inquiry into Small Forces [00:00:00]

Professor Michael McBride: Okay, let’s start up. So last time we got to the point of asking the question, are there atoms and molecules — and well, I anticipate that the answer is yes — and what force holds the atoms together? Because if we really understand the force that holds them, we’ve got chemistry all solved. So when you come to the idea of forces, in the first place, one goes back to Newton and Principia Mathematica. And he had the idea, or developed the idea, the mathematics of gravity. And the interesting thing about it was there was no mechanism for it. There were no springs or strings that connected the things. It was action at a distance.

So there were other ideas. Like, Descartes had the idea that gravity resulted from blocked repulsion, that the universe was full of particles zipping around and bombarding things. So like the moon would be bombarded by particles from all directions, but there’d be no net force on it because they’d all balance out. However, if you put the earth in the way, there’d be a shadow, and now you’d have more things hitting the moon from the right than from the left. So it would seem to be attracted to the earth. So that’s a great mechanism. Now, it was important for Newton that the force be proportional to r2, to inverse r2, 1/r2. Does this work for that? If you pulled the earth back, how would you change things? What do you say?

Student: I don’t know. It doesn’t work for this.

Professor Michael McBride: It doesn’t work? Would it block more or fewer things if you pulled it back?

Students: Fewer.

Professor Michael McBride: How many fewer? Suppose you made it twice as far back, how many fewer would be blocked from hitting the moon?

Student: Wouldn’t it be r2?

Professor Michael McBride: It would be 1/r2. You can think about that. So that aspect of it would work. How about — you know gravity is proportional to the mass of the things. How does that figure? You can think about that yourselves, okay? But anyhow, there were other ideas. But Newton didn’t think about the mechanism, he thought about the law, that it was 1/r2 and proportional to the masses. sirodonia

Now in 1717, he published the Second Edition of Opticks. He published the First Edition in 1704, and the reason he published it in 1704 was that Robert Hooke died in 1703. He’d actually — Newton had done this work in the 1660s and ’70s, but he didn’t publish it until Hooke died because he didn’t want Hooke to get on his case, criticizing it. But in the Second Edition, in 1717, he put some additional material; you see, “with Additions. ” And the additional material, even though the book is about optics, the additional material is about chemistry. And it’s at the end, and posed not as hypotheses but as questions, because Baconian people, following Francis Bacon, weren’t supposed to make hypotheses, but they could ask questions. So here are some of the questions. Question thirty-one, the one that has most to do with chemistry:

“Have not the small Particles of Bodies certain Powers, Virtues, or Forces by which they act at a distance, not only on the Rays of Light to reflect, refract and inflect them, but also upon one another for producing a great part of the Phaenomena of Nature?” (And this attraction between the particles or repulsion) “For it’s well known that Bodies act upon one another by the Attractions of Gravity, Magnetism and Electricity; and these Instances shew the Tenor and Course of Nature, and make it not improbable but that there may be more attractive Powers than these. For Nature is very consonant and conformable to her self. How these Attractions may be perform’d, I do not here consider. What I call Attraction may be perform’d by impulse…”

Where did he get that idea? Where did he get the idea that there are impulses that can cause attraction?

Student: [Inaudible]

Professor Michael McBride: That was Descartes’ idea, that we just explained. “…or by some other means unknown to me. I use that Word here to signify only in general any Force by which Bodies tend toward one another, whatsoever be the Cause. For we must learn from the Phaenomena of Nature what Bodies attract one another, and what are the Laws and Properties of attraction, before we enquire the Cause by which the Attraction is perform’d.” (So get the law first so you can do mathematics on it, and then worry about how it works.)

“The Attractions of Gravity, Magnetism and Electricity, react to very sensible distances…” (He doesn’t mean reasonable, he means ones that you can sense.) “…and so have been observed by vulgar Eyes.” (You can see the distances over which electrostatic attraction works, obviously gravity, magnetism.) “And there may be others which reach to so small distances as hitherto to escape observation; and perhaps electrical Attraction may reach to such small distances, even without being excited by Friction.” (So you know that friction can generate static electricity that’ll make things attract, but maybe even without friction you can get things that’ll attract a very short distance by electricity.) “The Parts of all homogeneal hard Bodies which fully touch one another, stick together very strongly.” Did you ever have that experience of having two really flat things, put them together and they’re very hard to get apart; what?

Student: Slides.

Student: Small things.

Professor Michael McBride: Microscope slides. Ever take out new microscope slides? They can be very hard to get apart. “And for explaining how this may be, some have invented hooked Atoms, which is begging the Question; and others tell us that Bodies are glued together by rest, that is, by an occult Quality, or rather by nothing, and others that they stick together by conspiring motions, that is by relative rest among themselves.” (These are other theories that are sort of complicated; we don’t need to go into it.) “I had rather infer from their Cohesion, that their Particles attract one another by some Force, which in immediate Contact is exceeding strong, at small distances performs the chymical Operations above mention’d, and reaches not far from the Particles with any sensible Effect.” (So a very short-range, a very strong attraction.) “…reaches not far from the Particles with any sensible Effect.”

So maybe it’s even more dramatic than 1/r2.1/r2 changes very rapidly when r gets very small, but maybe it’s even a law that’s more distance dependant than that, but only works at really, really small distances. “The Attraction” (and he’s talking about between glass plates separated by a thin film of the Oil of Oranges; so you know, microscope slides, they put a little drop of something and stick them together and they’re really hard to get apart) “may be proportionally greater” (the attraction) “and continue to increase until the thickness do not exceed that of a single Particle of Oil.” (Imagine Newton, 350 years ago or 300 years ago here, thinking about a single particle of oil. Could he possibly measure that? We’ll think about that later. Okay.) “There are therefore Agents in Nature able to make the Particles of Bodies stick together by very strong Attractions. And it is the business of experimental Philosophy to find them out.” That’s what we need to find, what is it that holds these particles together? “Experimental philosophy” means science. And we’re going to engage in this business, and it’ll take us about five weeks to figure out what it is that holds these things together.

Chapter 2. Is There a Chemical Force Law? [00:09:16]

Now, let’s just look, from physics, at the amount of binding energy you get from different kinds of laws. First there’s Coulomb’s law; so it charges over the distance. I think you all know that one. Then there’s magnetic interaction, two magnetic dipoles, and the energy is 1/r3, falls off very rapidly. Then there’s the “strong” binding that holds particles together in the nuclei of atoms. And gravitation, of course. And the chemical bond. So here’s a scale of energies, and we use, in this course, the old fashioned kilocalories/mol, which American organic chemists like to use. Some day it’ll change to kilojoules/mol, but for me it’s still kilocalories/mol. Okay, this is a logarithmic scale, so it covers a very wide range.

Now Coulomb, if you have a proton and an electron, as far apart as the distance of the carbon-carbon bond, that’s 216 kilocalories/mol, right in the middle of this, or near the middle. Magnetic interaction is much weaker. If you have two electron spins, at the same distance from one another, their energy is what? One, two, three, four, five orders of magnitude smaller, 105th smaller. The “strong” binding that holds a proton and a neutron together is in the deuterium nucleus, that is within the nucleus, not at 1.54 Ångstroms, is 105th stronger than Coulombic interaction. And gravity is ever so much weaker. If you had two carbon atoms, that were at that distance, the energy would be 10-32 kilocalories/mol. So forget gravity as a way of holding things together. But then we have Coulombic up or down by 5 orders of magnitude.

Now what is a chemical bond; how strong is it? A chemical bond is about — the carbon-carbon bond is about 90 kilocalories/mol. So not so far from the Coulomb energy, from electrostatic attraction. So maybe this hints that electrostatic attraction has something to do with it. But maybe not, maybe it’s something entirely new. And we haven’t talked about kinetic energy. Does that have anything to do with the strength of the bond? So we’ll get back to that later on. Okay, so is there a chemical force law, the thing that Newton was looking for? Now here’s an interesting way to phrase that. Suppose you had a chain of atoms, a thread that was single atoms bonded to one another, and you started stretching it. It would stretch, and at some point it would pop.

Now let’s just take three of the atoms and we’ll put stiff rods on the outer ones and we’ll start stretching it. So it’ll stretch, stretch, stretch, stretch, pop. Now the question is how far do you stretch it before it pops? What determines how far you stretch it before it pops? Okay, this must have to do with the force law between atoms, which then will have to do with molecular structure. So it could be like springs, Hooke’s law, Ut tensio, sic vis. Okay, so the force is proportional to the displacement, and the energy to the square of the displacement and the energy then, the potential energy, is a parabola. We define zero when they’re at the standard, unstretched distance, and then you stretch it or compress it and the energy goes up, like a parabola. Now the slope of that parabola tells you what?

Student: The force.

Professor Michael McBride: The force. That’s what’s linear. As we stretch it further and further, the slope gets bigger. The force is proportional to the displacement. Okay, now there could be other kinds of force laws, like electrical charges, Coulomb, or gravity. So that’s a different force law, inverse square force law, and the energy then is proportional to the displacement, in this way. So very different. Notice the zero of this one’s at the top when — zero defined when they’re very far apart. Here it’s zero when they’re at the standard distance. And so as you bring oppositely charged particles together, the energy goes down infinitely, as they get really close. Okay, and now if we look at the force here, as we go out the force gets smaller, the slope gets smaller.

Now suppose you had a particle that was connected to a spring on each side, as in this chain that we’re stretching. And just to be general I made the second string — oh I say stronger, it’s in fact weaker; it doesn’t go up as rapidly — I’ll have to change that. Okay now, but at some point the slopes of the red and the blue, at some distance, some position, will be equal and opposite. There. Okay, what’s the force on the central particle at that point?

Students: Zero.

Professor Michael McBride: Zero. They balance. So if we added those two energies together to get the sum of the two, on this central particle, it looks like that. So there’s a balanced minimum. There’s a well defined position where the particle in the center of this chain will just sit there and it’ll be hard to displace. So it can vibrate. Okay? Now how about in the other case? How about if it were electrical charges or gravity or something like that, an inverse square law that’s holding things together? Then the second one is going to look like that, and that one is indeed stronger.

So they’re two flanking bodies and we’re interested in the position of the one in the middle. And again, there’ll be some point where they balance, right there, where they’re equal and opposite slopes. How will this differ from the first case? Can you see what’s going to happen when we add these two together? Is it going to look more or less like the one on the left?

Student: [Inaudible]

Professor Michael McBride: It’s going to look like that. And now there is no balanced minimum in the middle. There’s not a position where the thing will just sit there, because it’s always more strongly affected, and attracted, by the one it’s closer to. There’s some place where there’s zero force but it’s not a stable position. Because if it displaces ever so slightly, it’ll keep on going. So that one is a — for the two of them — is a single minimum, and this one is a double minimum. And for the people who came early, there was a contest that if you won it, you could get an A in the course, and the contest is this. Here’s a magnet, hanging here. And I’ll stop it. Okay, and here are two magnets, the ones that will flank that one. And if I get it just right, it’ll just sit there, the attraction to the two will balance it. So I’ll put it on here, and try to get my A. So I’ll get the string where it would balance right there. Ah. Okay? There is no stable position. Nobody got their A that way. Right? Because it’s an inverse law, so it’s always more strongly attracted to the closer thing. So you can’t win with that, you can’t get a balanced position. Now I’m going to get these off here. And back to the show.

Chapter 3. The Morse Potential [00:18:27]

So with springs you might be able to make a stable polyatomic molecule from point atoms — we saw a spring model before in the last lecture — but you can’t do it with ions and you can’t do it with magnets. However, Hooke’s law can’t do it — it can’t be springs, because Hooke’s law never breaks. So we need a different kind of force law. Does anybody know a force law — what it looks like — the form of the force between atoms or the energy for stretching a bond? You ever seen a picture of such a thing?

Student: [Inaudible]

Professor Michael McBride: It’s called the Morse potential, or one form of it at least is called the Morse potential. And it’s not something fundamental, it’s not a law of nature. It was thought up by a physicist at Princeton in 1929, because it’s mathematically convenient; so you could solve quantum mechanical problems if you used Morse potentials. And the idea is this. You have two neighbors and you look at the position of the one on the right and it could be there and the energy would be minimum; that’s the bond distance, right? And if you move it further to the right, the energy goes up; move it to the left, the energy goes up. Okay, that’s the Morse potential, and we’re just holding the neighbor on the left fixed. Now we could put another neighbor on the other side and have another curve of the same sort; stretching the second bond. Or we could have it be a chain and have both neighbors there, and the sum of those two now has a minimum in the middle.

It’s a single minimum, like we got with Hooke’s law. So that means that atom in the middle would just stay there. And now we’re going to grab those sticks on either side and pull the neighbors apart. So if we pull them apart, it’s still a single minimum. And if we pull them apart further, it’s still a single minimum, although it’s very flat. And if we pull it further, it pops and it becomes a double minimum. And if we keep going the double minimum gets more pronounced. Now here’s an interesting question. At what point — how far do you have to stretch it before it pops? That’s the question we were asking, right? How can you look at those curves and tell how far apart you have to stretch it before the chain pops? Shai?

Student: The bottom is flat. So when the curve and the plot was flat, then that’s as far as it could go because…

Professor Michael McBride: What do you mean flat? There’s a name for a thing like that in graphs.

Student: Inflection.

Professor Michael McBride: It’s the inflection point. It’s when the curvature changes from being this way, like Hooke’s law, to being this way, right? So when the inflection points cross, then the chain pops. Okay, so force laws are going to make a lot of difference in how atoms behave, and if really knew the force law we’d be in a great position for understanding chemistry. It snaps at the inflection point where it changes from a direct to an inverse dependence on distance for force.

Chapter 4. What Are Bonds? Early Understandings of Valence [00:21:43]

Okay, but what are bonds? Newton said, “We’ll look for the law.” And it’s sort of, the law is sort of like Morse, but we don’t know where it came from, because we don’t know how it works, we don’t know what it really is. So can we find out what it really is? Well in the nineteenth century they discovered bonds. And this is a picture from 1861, we’ll talk about this later. It’s one of the first depictions that’s recognizable by people today of bonds between atoms. So there are different numbers of these lines, valences, for different atoms. Hydrogen has one; carbon has four; oxygen two; nitrogen three.

Why do the elements differ? Why don’t they all have the same valence? And even more complicated, sometimes nitrogen is three and sometimes it’s five. For example, you have NH3 but you also have NH4Cl, where there are five things associated with the nitrogen. So how do you understand the valence? That was the challenge for people. They had figured out that there was valence, but why? How could you predict the valence of a new element? Well Gertrude and Robert Robinson published a paper in 1917 which showed this picture of ammonia, NH3. What’s the loop? Pardon me?

[Students speak over one another]

Professor Michael McBride: Someone said — hold your hand up so that people can hear you. Okay yes.

Student: A lone pair on nitrogen.

Professor Michael McBride: A lone pair on nitrogen — wrong. It looks like the lone pair on nitrogen but the lone pair didn’t come along that soon. What the loop means, if you look at it in the context, is what makes NH3 reactive? In this term, this is what they wrote: “It may reasonably be assumed that the partial dissociation is a stage in the complete process.” So the bonds begin to break. This loop begins to break to become a weaker loop and two other partial valences, in this theory. And those valences begin to associate with partial valences from HCl. Right? The loop is a “latent” valence. You know what ‘latent’ means? Latent is the opposite of patent. Latent means hidden; patent is revealed. When you get a patent on something it means that you tell everybody how to do it but the government protects your right to do it for a certain period of time. Latent means it’s hidden.

So there’s a hidden valence, these loops, but they can become available, right? It must be assumed in some cases, as for example the combination of ammonia with hydrochloric acid. Now, so you can get a reaction and get the product, which has all five valences of nitrogen now. But how do you know there’s not another loop on nitrogen, that you could break open and make it seven valent, or nine, or eleven? Or maybe you could break a loop into three. So might latent valence loop explain the trivalence and pentavalence of nitrogen, or the amine-HCl reactivity? The trouble is it’s too slippery a concept. It explains everything. You could explain anything you wanted to. You could explain eighty-four valence of nitrogen or something. Anything that comes along, you’d say, “Ah ha, there are that many latent valences.” But how do you know there are not more? Why do you have latent valences? When do you have them? When and why do you have partial dissociation? This thing didn’t explain anything, or explained everything. Now, at this point I want you to smile so I can learn who you are.

[Professor McBride takes pictures of the class members]

Professor Michael McBride: So the electron was discovered in 1897. So maybe that has something to do with it. And the guy who had associated electrons with valence was G. N. Lewis. And this is a picture of him as a Harvard undergraduate in 1894. And this is eight years later, when he was an instructor at Harvard, and he then went to establish the Chemistry Department at the University of California at Berkley.

Student: Yay!


Professor Michael McBride: But when he was at Harvard in 1902 he used these lecture notes, and what he was trying to explain was the periodic table, why you go across in eight and then another eight and another eight, and so on, and why you have electron transfers from some atoms to other atoms. And he said if the electrons in an atom are arranged at the corners of a cube, then you — eight is this very special number, because that occupies all the corners of the cube. So if there’s a desire, for some reason, to complete octets, then you can explain periodicity and see why some atoms give up an electron to lose the outer shell and others gain one to complete the outer shell.

Okay, so but it also could explain how you get bonding, if you like to have octets; not just electron transfer but the formation of covalent bonds. So, for example, here’s two chlorines. Each has only seven; the octet isn’t complete. But if you bring them together and they share an edge, then both octets are complete. He doesn’t say why octets should be so great, but if they’re octets then you could explain bonds. How about double bonds? Suppose you had two oxygen atoms. What do you do now? Anybody see? What?

Student: Put the faces together.

Professor Michael McBride: Put faces together, share two edges, and then both have an octet. Now, suppose you want to put nitrogens together and form a triple bond. Now how do you do it?

Student: Squish it.

Professor Michael McBride: Push it hard, force it like a picture puzzle?

Student: Yes.

Professor Michael McBride: Put it in there.

Student: You need to push it.

Student: You did it.

Professor Michael McBride: It won’t do it. But about ten years later he figured out how to do it. What you do is take the electrons — nitrogen has 5 — but instead of using an octet, you contract it along opposing edges this way. So what was an octet, what was a cube, becomes what?

Student: A tetrahedron.

Professor Michael McBride: A tetrahedron. And now if you take two tetrahedra, you can put a face together and complete both octets. But this wasn’t such a great development. He did this, published this, in 1916. But the tetrahedral distribution of the bonds in carbon had been known by organic chemists for 40 years by that time, that the tetrahedron was a very important structure. So a good theory should be realistic and it should be simple. But there’s a tension between these things. It should at least be as simple as possible, but that may not be very simple; that’s why it takes us five weeks. In regard to the facts it should allow a number of properties. It should allow prediction. You should be able to say how will these atoms react; how will this molecule react; what valence should nitrogen have; how can it be both pentavalent and trivalent? A little below this on the scale of desirability is suggestion. A theory may — even if it doesn’t give you the proper thing, may at least suggest some experiment that would be really good to do.

Below that is explanation — according to my theory all these known facts fit together, but who knows about any new thing. And lowest is classification and remembering; like Roy G. Biv allows you to remember the colors of the spectrum, but there’s nothing fundamental about that. So even if you have a crummy thing it might allow you to at least remember some facts and organize them maybe. Okay, but those last two are not prediction, they’re post-diction; you only use it to explain things that are already known as a way of remembering them. So from the number of valence electrons we’d like to predict how many atoms of different types come together to form molecules. That’s constitution — the nature and the sequence of bonds — and the valence.

Chapter 5. Deriving Structure and Reactivity from Valence Electrons [00:32:52]

We’d like to know the structure of molecules, the distance between atoms, angles. We’d like to know how charge is distributed within molecules. Are some regions positive and others negative? We’d like to know what the energy of the molecule is. Is this arrangement of the atoms better or worse than an alternative arrangement? And we’d like to know about reactivity. So Lewis explains constitution, the nature and the sequence of bonds. “Nature” means like single, double, triple, and “sequence” means which atoms are connected to which. So he has electrons, valence, and also he added unshared pairs. So you count up how many, from the atomic weight or the atomic number or the position in the periodic table. You get how many valence electrons, how many of these outer octets are there. And then you can figure that the valence of hydrogen should be one, boron three, carbon four, but then down again, three, two, one.

So here’s ammonia, NH3. It works fine and it has an unshared pair. But why do you have octets? Why not sextets? And why do you have a pair for hydrogen? Why not an octet for hydrogen also? Okay, so let’s just apply it to HCN. So if carbon has four, that makes possible a single bond to hydrogen, a triple bond to carbon. Everybody has their octet except hydrogen that only wants two. So everybody’s happy. And we can abbreviate it this way, and leave the dot. Lewis actually thought of drawing a colon, a pair of dots, to denote an unshared pair. So that’s his notation. So NH3. Bingo! How about BH3? Just fine. But here’s something new. The organic chemists would already have known this, they knew the valence numbers. The electrons just added something to it but didn’t say anything really new. But here’s something new, that BH3 reacts with NH3, to give a new bond. Now here’s a puzzle. It’s also true that BH3 reacts with BH3. That doesn’t look like Lewis. And the bond is reasonably strong, it’s half as strong as a carbon-carbon bond. So what’s the glue that holds B2H6 together? And we’ll come back to that fourteen lectures from now.

But there’s another thing. There’s bookkeeping that you do with Lewis structures to assign what’s called formal charges; not real charges, just what charges you write in the formula, and you hope they mean something. And the bookkeeping scheme you use is that each atom is assigned half interest in the bonding pair. So if there’s a bonding pair, for bookkeeping purposes you assign one electron to each. It may be they aren’t evenly shared, but the bookkeepers don’t care.

Incidentally, speaking of bookkeepers not caring, there’s — people have had a lot of nonsense trying to get signed up for lab, because the people who run the system don’t allow teachers to have access to the thing to see the problems you face. So I have no way of telling — I can’t get on to your system. Has everybody figured out how to sign up for lab by now?

Student: Yes.

Professor Michael McBride: If anybody still has problems, after class Dr. DiMeglio can take you to the computer center here and help you. So next year presumably it’ll be easier. This is a new system this year. But anyhow, back to the bookkeepers here.

So these formal charges are not necessarily real charges, they’re just what you write when you draw the formulas. Okay, so we split the BN electrons between the B and the N, and we split the hydrogen bonds as well. So nitrogen now has those four electrons, where it brought five to the table originally. So it has a positive charge. Tetravalent nitrogen is positive in the Lewis structure. Why? Because in forming a fourth bond, it gave up half-interest in one of its electron [pairs], the ones in the unshared pair. So it lost a charge, a negative charge, and becomes positive. By the same token, what else can you write in this structure?

[Students speak over one another]

Student: It’s negative.

Professor Michael McBride: A negative charge on boron, because it got half-interest in a pair. But that’s just bookkeeping. So when you write a Lewis structure you’d write plus on nitrogen and minus on boron. Now is it real? Is there a positive charge on nitrogen and negative charge on boron? Well we’ll get later on to know about what a surface potential is. Surface potential is the energy a proton would have if it was on different points on the molecular surface. For this purpose you have to define what the surface is of BNH6. There’s a way of doing that; we won’t talk about it now, that’s jumping the gun. But that’s the surface, and it’s color coded to show the energy a proton would have if it were on the surface at that point. So what you see is that the energy of the proton would be high, if it’s here, where it’s very blue. The energy of the proton would be low if it’s in this region here. Is that consistent with what we were talking about?

Student: Yes.

Professor Michael McBride: Where would a proton be low in energy? Where there’s negative charge, and it would be high in energy when there is positive charge. So this molecule, according to quantum mechanics, is positive on the left and negative on the right. And that’s exactly what we have, it’s consistent with what we have. And now Lewis also explains where pentavalent nitrogen comes from, because there’s an ionic bond there too. Tetravalent nitrogen is plus. So there’s just a Coulombic attraction between ammonium and chloride. So that’s a great contribution. Now there are other things you can understand in terms of Lewis structures. Like you have an amine and it can react with an oxygen to give an amine oxide, which is positive on nitrogen, negative on oxygen. Or you can have a sulfide, which can react with an oxygen, to give a sulfoxide. Or what else can happen, with a sulfur, what additionally can happen?

Student: Another oxygen.

Professor Michael McBride: It could do another oxygen, and give a sulfone, as it’s called. So Lewis explains this. So there are problems to drill on Lewis structures. They’re on the Web, go look at them. One of them is draw Lewis dot structures for HNC, in that order, H then N — we did HCN; but do HNC. And then try HCNO with CNO in all six linear orders; that is, nitrogen in the middle, carbon in the middle, oxygen in the middle, and then hydrogen on either end. And it’s also possible to make a ring of CNO and put the hydrogen on any of the positions. So just practice drawing Lewis structures of those.

And start memorizing functional groups. We do precious little memorizing in this course, compared to most organic chemistry courses. We try to figure out how things work so that you can predict without memorizing it. But there are some things you have to know, like the names of groups. Otherwise it takes forever to talk about them. So here are the ones you have to learn, and on the fist exam, coming up in three weeks or whatever, I’ll give you a question where you have to draw one of these structures from the name, or give the name for a structure.

These are all on the webpage; you don’t have to copy them down. But those are simple functional groups. [Laughter] And here are — whoops, oh there you go, sorry. [Laughter] There. They’re also conjugated functional groups. So you could run it yourself. Okay, now the final question here, for today, is a thing that comes up when you talk about Lewis structures, or structures in general, Lewis structures in particular, is equilibrium and resonance. How many people know what resonance is? What is it?

Student: Like the same molecule but different structure.

Professor Michael McBride: Speak up a little bit.

Student: It’s the same molecule but a different structure.

Professor Michael McBride: The same molecule but a different structure. But isn’t structure a molecule?

Student: But with that I guess —

Professor Michael McBride: You know like, let me just interject a minute here. Like, one of the problems is HCNO in any order. So there would be different molecules with C in the middle, N in the middle, O in the middle. Is that what you mean?

Student: No. I mean like it’s — if you have a double or a triple bond, it’s in a different location, it’s around the central atom.

Professor Michael McBride: Ah, now let’s — actually resonance is one of those things where people try to hide ignorance from ignominy. And let me show you what I mean by that. So here’s the structure for HCNO. And here it has all octets. If you go through you’ll figure that out. It has charge separation though. You don’t like to have that in your structures. You like not to have charges if you can avoid it. Obviously in BNH6 you can’t avoid it. Okay, now but you can get rid of the charge separation by shifting that electron pair that’s on oxygen to be shared between oxygen and nitrogen; then you don’t have the charge on oxygen anymore. Everybody see that?

Student: Yeah sure.

Professor Michael McBride: What’s the problem when you do that?

Nitrogen now has too many electrons. So you could shift a pair from — that’s in the CN bond, onto the carbon, in your drawing you make. Okay, so you get this structure shown below which has all octets. But there’s still charge separation, and in fact it’s worse because the negative charge, instead of being on oxygen, is now on carbon. So there are these two different Lewis structures you can draw and each has its own properties, or good things and bad things.

Now what are the geometric implications of there being two structures? It could mean that in the middle one, in the lower one, nitrogen is exactly halfway between carbon and oxygen, because they are both double bonds. But in the top one, nitrogen is much closer to carbon than to oxygen. That could be the meaning of this. So that if you draw a picture of the energy, as a function of the position of the nitrogen, you could get two different structures, one where it’s further to the left, where the nitrogen is further to the left you have this structure, and when the nitrogen’s in the middle you have that structure. And you have to go up hill in energy, it’ll click; as you start pushing the nitrogen it starts, it’ll click, and it’s the other one. Okay? That could be the way the atoms really behave. That’s called a double minimum; we already saw a double minimum today.

But maybe in truth it’s not this way. Maybe it’s a single minimum. Maybe the best position for the nitrogen is neither to the left nor in the middle, but in between; maybe that’s the lowest energy. And that is what resonance is. Resonance is when you — the true structure is in between the structures you draw. So really what it is, is a failure of the notation. Does everybody see how that’s so? It’s that the notation you use doesn’t show the right structure, and when that happens there’s said to be “resonance”; although actually all it is is that the true structure is a single minimum, not the double minimum you’d expect from your drawing. It’s a failure of the way you draw.

And notice that for resonance you use a double headed arrow whereas if I back up a second here — equilibrium, when you have a double minimum, there are two arrows, two different structures. So the question is which is it? Is the real molecule a double minimum; or the real molecules a double minimum, are there two structures? Or is there in truth only one structure? So you need to choose. The choice between whether there’s resonance or equilibrium must be based on experimental facts or on a theory that’s better than Lewis’s theory, to be able to know which is true; something that can distinguish between a single and a double minimum.

So equilibrium is two arrows, and there are two structures that go back and forth, and resonance is a single structure, that our notation isn’t very good at drawing. So two real species, or one real species but two reasonable formulas we could draw for it, or that Lewis would draw for it. So really what it is, resonance is a failure of simplistic notation, and it’s associated with unusual stability. Now when I say ‘unusual,’ that implies that you know what usual stability is. So compared to what is it unusual? And we’ll address that question later on.

So let me just give you finally an example of equilibrium versus resonance. So this, as you’ll learn when you learn functional groups, is a carboxylic acid, and you could have the hydrogen attached to this oxygen or it could be attached to the other oxygen. So these are not just the same thing upside down. It’s the hydrogen being attached to one or being attached to the other of the two oxygens. And, in fact, it is that way. There are two structures and you go back and forth between them; two species.

But how about if you take the hydrogen atom away and have this thing that’s called a free radical, where you don’t have a complete octet on this oxygen? There are only two, four, six, seven electrons there. It has the ability to form another bond by sharing an electron. Is that two different species? Or is it one species — watch — with an intermediate geometry where you don’t really have double and single bonds but something in between? And how do you know which it is? The only way is to do some experiment that tells you. And there is an experiment that tells you it’s only one nuclear geometry. And we don’t have time in the last thirty seconds to tell you how that proof works, but I’ll tell you that the technique is electron paramagnetic resonance spectroscopy [Laughter], shows that in truth, although this is two species, this is one species. So there’s equilibrium at the top and resonance at the bottom. And if you add an extra electron to make a carboxylate anion, infrared spectroscopy tells you that it’s symmetrical too. There’s resonance in this case. It’s an intermediate structure. But that is “lore.” You didn’t predict that from Lewis structure ahead of time. It could’ve been either way, and in fact there were many people who thought it was one way or the other. It’s only experiment that does it. So, next time we’ll talk more about lore.

[end of transcript]

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