We'll also take a look at the end at a few more specialized examples and see some things that you can tell with IR and just a few neat things and ways it can talk to you. But let's start with some basic nitrogen containing compounds. So amides continue our discussion of carbonyl groups.

We talked about the amide carbonyl band and we said of all the common functional groups, of all the common carbonyl compounds, amides have the lowest carbonyl stretching frequency at 1650 to 1690 and sometimes this is referred to as the amide 1 band

and we'll see in a moment what the amide 2 band is. Now what's interesting is you often end up looking for things in pairs, looking for sets of data that point

in a particular direction with IR spectroscopy. So the other thing for primary and secondary amides, that is amides where you have either two hydrogens or one hydrogen on the nitrogen is you're going to have NH stretching frequencies. So the primary amides, I'll write it as primary amide,

compounds where you have NH2 groups have two NH stretches and by now you should all know what are those NH stretches in terms of their type?

Symmetric and asymmetric and their frequency is going to depend, obviously it's going to be, I think everyone by now knows that that region above 3000 is

where you get OHs and NHs and so forth and it's going to fall in that general region. Amides have, just like carboxylic acids, they're very prone to hydrogen bonding and so often you'll see different behavior with different states. So for a primary amide in solution and I'll say dilute,

in other words not super concentrated so you don't have a lot of hydrogen bonding, those bands show up at about 3520 and about 3400.

Remember my emphasis really has been on reading spectra, in other words looking at a spectrum and seeing patterns here and so when I look I sort of look at that region below 3000 and I'd be seeing a band

and these bands tend to be kind of broadish, not as broad as your typical alcohol bands but certainly not as sharp as a CH band. Hydrogen bonding weakens the NH stretch and so

in the solid state they're certainly hydrogen bonded even when you grind up your molecules with KBR or grind them with Nujol, grind them with mineral oil to make a mole, the molecules are still in particles, micron and size or so forth so they're all hydrogen bonded together.

So in solid, I'll put parenthesis hydrogen bonded, the bands shift to lower frequencies and they typically broaden out. I'll say about 3550, about 3350 and 3180 and the pattern changes a little bit, I'll try to graph things

on the same scale here so now we fall below 3500, the bands typically become a little bit broader so you'll see this kind of like an alcohol stretch

over here becoming broader. So just like we talked about carboxylic acids, amides particularly primary amides and lactams can form these hydrogen bonded ring structures and are quite prone to doing so and the reason,

so primary amides are more prone to doing this and as are lactams and the reason is that by the time you get to a secondary amide, that is where you have one R group on the nitrogen,

typically the hydrogen prefers to be trans to the carbonyl group so it doesn't set itself up as nicely for forming ring structures. You can still form chains but they don't tend to form as extensively.

Secondary amides also have an N8 stretch so again,

what I drew over on the other blackboard is a secondary amide and typically you're talking about an N8 stretch from about 3400 to about 3500

and if you're hydrogen bonded or about 3300 if it's hydrogen bonded. So in addition to these bands which are called the N8 stretch bands, as I said,

you also have for primary and secondary amides, what's called the amide 2 band. The amide 2 band is an NH band and that typically is

on the order of about 1500 to about 1650 and it may fall on top of the amide 1 band, in which case you wouldn't see it.

All right, I'd like to help us get in the habit

of the pattern recognition on these and start to see the differences in these classes of compounds just so you get good at seeing what's out there. So let me pass out, I've got some handouts here. Let me send these down and we'll look at a few spectra.

I might need one more.

Send the extras back there.

All right, so this might also be a nice time for a little bit of a question here.

So we're on the subject of amides, so we know we have an amide here. Problem here, this was taken from a book,

problem says we have a compound with a molecular formula C6H13NO and it contains the tert-butyl group. So can we figure out the structure of this compound?

Anyone have an idea?

So a methylene group that's attached to the tert-butyl group.

So we've used up, used up our, so what do we do with this?

So we have two spectra here. We have one spectra in chloroform solution and I'll tell you it's a concentrated spectrum. So this is, you notice these numbers here.

So we have one number at 3400 and another number at 3300. So clearly, clearly we have at least one hydrogen bonded, at least something in our NH's is hydrogen bonded and they don't look quite like that. So is it primary or is it second?

It's got to be a secondary amide. So what do we do with the structure? So a CH3 and a tert-butyl and like so or the other way

and we, with this level of information, we can't tell the difference between it and in fact, in fact this structure is the first one here.

Now, there's one question. Ah, okay. So we have, so this is tert-butyl acetamide. It's a secondary amide. This is our free NH.

This is hydrogen bonded. So typically you'll be working with like a 5 or 10% solution. So many of the molecules are monomeric and you see a band for the NH stretch at about 3400.

You're also going to have some cases in which the molecules are hydrogen bonded together like so, either to form two or more molecules and so the hydrogen bonded NH stretch is going to be the one

at about 3300 wave numbers. For a primary amide, they're a little bit different in size. Usually the one that's the higher wave numbers is slightly longer in length.

Ah, okay. So great question. So what experiment could we do with this sample in chloroform to determine whether the band at 3300 was associated with hydrogen bonded dimer formation?

Diluted and what would you expect to see? The decrease in the 3300 or more specifically, both bands would of course decrease because it was diluted but the relative intensity of the band at 3300 would get low.

Based on the spectrum alone, if I looked at this pattern here, I'd say well this clearly doesn't look like a monomeric primary amide because I said you'd expect at about 3520 and about 3400 and it probably doesn't look

like a hydrogen bonded primary amide. First of all, this band at 3300 is a little bit high and it's a little bit narrow. Remember I said typically 3180 but if you saw my sketch, it was sort of broad like an alcohol where it basically spanned about 300 wave numbers.

So you might or might not be able to do this with pattern recognition. I mean this is as you see more of these spectra you'll get better and better. Now what's happening here in the solid state is you're only having the hydrogen bonded band so you're seeing just the hydrogen bonded NH.

Other thoughts and questions? These are good ones. The little shoulder right here.

The peak, which spectrum? The peak at the shoulder that's on there. Probably different states of hydrogen bonding. Probably some with more hydrogen bonds. For example, you can have a bifurcated hydrogen bond.

So it may be something like that. It's something clearly that's characteristic of the solid state. Now we're dealing with, and again this is what I was saying

for this sort of pedagogy of spectroscopy. We're dealing with a relatively small molecule here. It's only got six carbon atoms. By the time you get to bigger molecules, the relative intensity of your NHs relative to, for example, your CHs is going to be less. And this is where I said KBR is really bad

because you typically, it's hard to get out every last bit of water and it doesn't matter so much for a molecule this size. But if you're dealing with a molecule say the size of strychnine, you're going to have almost 10 times, let's say five times the carbon to nitrogen ratio

or carbon to amide ratio. So that peak's going to be a lot smaller. You're really going to have trouble seeing it. Okay, the one other thing I wanted to bring up was just about problem solving strategies on this. And I presume everyone's seen this, but it's worth bringing up the idea.

So this molecule was C6H13NO. And the first thing that you want to think about when you're asking what's this compound is the unsaturation number. And you can do that mathematically. In other words, how many double bonds and rings are in there?

That's an incredibly useful piece of information. In fact, molecular formula is probably the most useful information that you can get for starting with a compound followed by what functional groups are in it, which is why we're starting the course with IR and mass spec. So there are a couple of ways you can do it. Honestly, I never do it the way I'm about to write it.

The unsaturation number is the number of carbons minus the number of hydrogens over 2 minus the number of halogens over 2. I hate formulas. I don't even keep formulas in my head. Plus the number of nitrogens over 2 plus 1. And if you plug in, right, that's 6 minus 13

over 2 plus 1 over 2 plus 1 is equal to 1, which is, you know, one double bond, a carbonyl group. I never do things this way. Basically, the way I think about it is if a molecule has 6 carbons in it, if it were an alkane,

it would be C6H14. If it has halogens in it, they just count as hydrogen. So in other words, C6H14 is saturated, C6H13CL is saturated. Oxygens don't count toward unsaturation number. Sulfurs don't count toward unsaturation number.

Nitrogens, you need one more atom in there, one more hydrogen in there to complete the valence. In other words, an alkane would have a formula C6H14. If I go ahead and add a nitrogen, it would have to be C6H15.

And then I simply look and say, okay, it's H13. So there's one degree of unsaturation. So that's how I think about unsaturation number. But then immediately you're saying, wait a second, we've got a carbonyl in here. It can't be a lactam. It can't be a ring compound, for example.

We only have one degree of unsaturation. Thoughts or questions at this point? Why is the carbonyl lower here? The hydrogen bonding. So the hydrogen bonding ends up weakening the carbonyl.

It's kind of, what's kind of counterintuitive on hydrogen bonding is you can think of hydrogen bonding as being an electrostatic phenomenon primarily. And by that I mean you think about your delta positive

on the nitrogen and your delta negative on the oxygen and you're pulling electrons away from the double bond. You're weakening the double bond in forming a hydrogen bond,

in forming that electrostatic interaction. Then you think about the NH stretching frequency. You have electrons on the oxygen. Those electrons are pushing the NH electrons toward the nitrogen. They're weakening the nitrogen hydrogen bond. So that stretching frequency drops to a lower number.

All right, so let's do another class of nitrogen-containing compounds.

And you notice for the amides, particularly in the solid state, the stretch is a little different than alcohols. I'll show you an alcohol in case you haven't seen it in just a moment, but alcohols tend to be less lumpy

than hydrogen bonded amides in the solid state. They tend to be pretty ugly. All right, let's take a look at amines. Amines are similar to amides, but weaker.

Obviously you don't have a carbonyl, but in other words you have NH stretches, you have NH bends, but the bonds aren't as heavily polarized. And when you have less polarization of the bond, you have less change in dipole moment, right? The carbonyl on the nitrogen is very electron withdrawing.

So that increases the change in dipole moment. It gives you a stronger band. So primary amines, RNH2, you'd see two bands are symmetric and asymmetric stretch.

The asymmetric typically is about 3500 free and about 3400 for the symmetric stretch. In other words, these are the, I'll just write free here, not hydrogen bonded.

Amines don't hydrogen bond nearly as much as amides or alcohols. So in solution you're not going to see them hydrogen bonded in the solid state or anneat, meaning the pure liquid. For example, sandwiched as a film between salt plates you will.

So anneat amine typically you'd have about 3400 and about 3300 secondary amines, R2NH. Now you don't have the coupled vibration.

So you see one band slightly lower in frequency. Let's say approximately 3350 to 3310.

So when you're looking at data like this, you're not necessarily going to have something scream to you this is definitely an amine. But what you're going to have, or this is definitely an amide, but you're going to be getting pieces of evidence and saying wait a second, the data is pointing in this direction, a red flag is going up.

This seems to be consistent with an amide because I'm seeing a carbonyl and an NH stretch or two NH stretches and then you start to look for other data or you perform additional experiments. Let me give you an example here on an amine.

And so I wanted to just put contrast here. So I've taken both an amine here

and I'll tell you what these are this time. This is butyl amine and this is butanol, one butyl amine, one amino butane and one butanol. And you'll see some differences. So here's what we see for a typical alcohol OH stretch.

They're both neat, the alcohol is hydrogen bonded. Strong, it's broad and again in terms of pattern recognition you sort of see it coming up at 3500

and going down by about 3000. The amine is weaker, in this case because it's neat and because it's primary we are seeing it hydrogen bonded. I'll tell you secondary amines are often really hard to see.

They're often very weak and by the time you get to a molecule that has more carbon and less amine in it they end up being darn hard to see. What's frustrating with secondary amines, primary amines too but secondary amines in particular is often in the NMR spectrum particularly

in chloroform solution. It's really hard to see that NH, the NH band in the NMR spectrum so it can be really vexing and if you're working say for your aurels and you know you have an amine because let's say you've carried out a reaction, everything else is consistent, the mass spec is consistent,

you'll end up saying oh my goodness I'm not seeing the NH resonance in the NMR, I'm not seeing an NH stretch in the IR and you might have to look really hard for the NMR you might want to dissolve it in DMSO which tends to be a good solvent that avoids exchange.

Thoughts or questions?

So let's take ammonium salts. If carboxylic acids are the ugliest thing you'll see, ammonium salts are even uglier.

If carboxylic acids are pugs, ammonium salts are bulldogs. So ammonium salts, the NH band is just huge and misshapen is really the only way to describe it

and again here's my very simple minded view of an IR spectrum. I'll put in a mark at 2000 here because I'll be drawing upon that in a second.

So for ammonium salts the NH, so I'm talking like RNH3 plus or of course you could have just like a secondary ammonium salt but basically

at about 3200 things come on and you know, you might see some CH's poking out and then they'll come down at 2000 to 2500 so here you start

to inflect at about 3200. Of course then you'll have whatever else you have in the spectrum but it's basically let's say about 3200 to 2000 or 2500, very broad, very ugly.

So one reason that this is so important to keep in mind is one very common laboratory operation is

to isolate an amine by liquid-liquid extraction. If you're teaching sophomore organic chemistry you're probably doing that, making a free base from a hydrochloride salt or taking an amine into an acid layer and then taking it back and it's very easy,

you know even in a research project to end up getting yourself fooled and think you made a free base, people will often use say sodium bicarb which is just the bicarbonate or the conjugate acid for bicarb is just matched to an ammonium group so it's not a great way to free base an amine.

You're much better off with sodium hydroxide or sodium carbonate but I've seen this happen to people where they get a product and they say my amine doesn't dissolve in anything. It doesn't dissolve in organic solvents. Why isn't it dissolving? I stripped it down, you know, I isolated it, I precipitated it and it's not dissolving. It's because it's an ammonium group

and that's something that would be hard to tell by NMR spectroscopy. If you knew exactly where to look you might say oh, the adjacent methylene is often chemical shift by a few tenths of a PPM but an IR spectrum where you see something like this it's like oh my God, you know, clearly my compound is talking to me.

Let me show you an example and this is a cool one. I pulled this last night. I don't think I included this in the handout so I think I need to give a separate handout.

We need to send a few more back here. Maybe we send these back. What's?

Great question. So the question that was just asked is if we're talking about a primary amine what's going to happen if we change the group here to like a hydroxyl amine? It should be pretty similar because to a first order approximation your reduced mass

is going to be the same for your NH bonds and your bond strength will be very similar so your root K over mu term in your harmonic oscillator is going to be similar. For specialized compounds like this, this is a great place to look things up.

If you really needed to know the exact frequency the book you have, the Prech book is really good for looking at specialized compounds. At the end of today's class I'll give you some general principles of where to look for things because there's a lot of overlap, in other words basically it doesn't change that much.

Heteroatom to hydrogen bond. We saw all of our alcohols, our carboxylic acids, our amides, our amines are all in that general region between about 3,000 and about 3,500 but if you needed something more subtle you could look that up. All right, so the example that I picked here is phenylalanine.

I seem to have, what am I going to do with my marker? So, all right, so phenylalanine if you look at your sophomore organic chemistry text you'll see

that's the structure of phenylalanine but you'll also see and not surprisingly since the pKa of a carboxylic acid is about 5 and the pKa of an ammonium group is about 10 that you have an equilibrium

that lies essentially completely to the right toward the zwitterion and it might be hard

to see this by some other method but you've got this really, really typically ugly pattern where you're starting at about 3,200, you're coming through this lumpy region here for all the different hydrogen bonded states and in this case coming down at about 2,000.

So it's really, really talking to us. Thoughts or questions?

All right, I want to take a couple of other nitrogen containing compounds and then maybe talk about a few other things.

All right, nitriles are carbon triple bond in nitriles. Those always surprise me because I expect the carbon nitrogen bond to be highly polarized. I expect it to be stronger and it really isn't.

The position ends up being very precise. It's about 2250 wave numbers. Doesn't vary a whole heck of a lot so it stands out. It's in that region between 2,000 and oh about 2,700 where you don't normally see a lot of stuff.

It's sharp and typically it's weak to moderate in intensity. In other words, if you've got a big molecule and you have a nitrile functionality in it, you'll see a little blip there and it'll be recognizable and sharp but it won't be a big band like a carbonyl.

Let's take another class of compounds. We'll take nitro compounds so RNO2 or aromatic nitro compounds. I'll just draw the nitro group.

So we typically see two bands associated with the nitro compound, associated with the NO stretch, about 1,500 to 1,600 and about 1,300.

They tend to be strong. The one at 1,500 to 1,600 tends to be a little stronger. What are the two bands? Symmetric, asymmetric.

Even though I'm writing one resonance structure that doesn't mean there's one nitrogen oxygen double bond and one nitrogen oxygen single bond. There are two equal resonance structures contributing both

equally and so it's one and a half bonds a piece. The asymmetric stretch is always higher in frequency. It's just a higher energy stretch. Same question.

All right, so by way of corollary, let's see what you can do with ideas here. So what do you think about carboxylates?

What do you think we'd see for those?

What? You mean lower frequency? Indeed it is and that's more single bond.

In fact it's one and a half bond character for exactly the same reason as nitro groups. And remember your reduced mass means you have a square root term so if you're dropping your bond strength

to one and a half times as less your frequency is not going to be one and a half times lower. It's going to be like 1.2 times lower. How many bands do you think you see for it? Two? Two or one? How many think two?

Two for exactly the same reason. So in other words you'll see about 1650 to about 1550 for the asymmetric stretch and about 1,400.

And again if you're thinking back to this issue of making a free base then if you're dealing with a carboxylic acid and you're protonating it and you're saying oh my carboxylic acid isn't dissolving in any solvent. It doesn't dissolve in chloroform solution. I can't get an NMR spectrum of it.

Why not? It might be that you have a carboxylate anion and even if you dissolve it in DMSO you probably are since often you don't see carboxylic acids in the NMR you're going to not be able to really tell by NMR spectroscopy but you take an IR spectrum and it's like well I know I should see a carbonyl

for a carboxylic acid at about 1,700 or just a hair above 1,700. I'm not seeing it. Oh yeah I see this funny band at about 1,600 and you say ah my carboxylic acid isn't protonated. So again this is a case where IR really, really can shine.

All right obviously we can't talk about every functional group known to man and if your research project focuses on a particular functional group, so if your research project for example were to focus on making isocyanides say in some

of the marine natural products you would probably go to Prech and look up where the isocyanide functional group showed up and then when you're carrying out the strategic reaction to introduce that group you'd be able to look in the IR. But at the same time for reading an IR spectrum you can

make some generalization. So here again is my little sketch. I'll just put in my marks at the key places I like to look. We've already talked about this region here for NH and OH.

The region here from about 2,000 to 2,300 is an unusual region. If you see anything other than a little bit of CO2 from breathing into the spectrometer you probably

should have your head perk up because you have many triple bonds in this region. So even though I don't know where an isonitrile function shows up I would expect it to be in that region.

You also have accumulated double bonds. Cumulated means where you have two double bonds connected to the same atom. We saw carbon dioxide showing up in that region. Another common functional group that's often used

in synthetic chemistry are allenes. Allenes show up in that region as well. I can't name every double bond but double bonds in general have about the same strength.

You're dealing with typically elements like carbon, nitrogen, oxygen, all of those elements will give you a similar reduced mass. So this region from 1,800 to 1,300 has many double bonds. In other words your carbonyl stretches are there but also

if you're doing a project with imines that would be the first place you'd expect to look for imines. If you're doing a project with nitroso groups that would be the first place I'd expect to look for nitroso groups. We've seen carbon-carbon double bonds in that region as well.

All right, I'm going to mention a couple

of more specialized pieces of information just as examples of things you can see with IR spectroscopy. So aromatics, you have combination bands involving the carbon ring at about 1650 to 2000.

Those bands are weak and they tend to be a little bit on the fat side and they often show up in patterns. You'll also, and you can be diagnostic by pairing things,

have out of plane bands, CH bands at about 675 to about 900

and I'll just show you patterns that have been documented. Is this, I think this is part, can you check whether this is the last page in your handout?

Two more pages, great. All right, so the next to the last page in the handout. So for example, phenyl groups are very common and you'll see this pattern of four lines here

and you'll also see, so this is your combination and this is your out of plane. This pattern, you'll see they look very different than a carbonyl stretch. They're much weaker. You'll see this come up in a couple of your homework problems in the four-page sheet.

One of the homework problems says identify the functional group present in the carbon, hydrogen, oxygen containing molecules. One of the molecules that you will encounter there is ethyl benzene or toluene, I think it's ethyl benzene and you will see this pattern there.

Another problem you'll have which will be an interesting one is on polymers and you'll see this pattern show up along with some other things in that problem on identifying polymers. So you can pick out different substitution patterns. You could for example pick out ortho xylene versus meta xylene

versus para xylene based on the patterns here. Let's see, a couple of other special topics here.

One thing that's really cool and again this will come back to that homework problem on polymers, so for CH3 groups there's a very, very sharp band at 1375 that's a CH bend, it's the CH symmetric bend

for methyl groups and remember how I talked about coupling. If you have two methyl groups together you get a coupled vibration so in an isopropyl group, in an isopropyl group

or a gem dimethyl group you will see two bands, one at 1370, one at 1390 due to coupled CH bends, symmetric

and asymmetric types of things. That will actually help you in determining some of your polymers apart. The last thing I'll just show you for the heck of it because it's on the last page of the handout

and then I think I'll wrap up on IR. Just one more sort of special topicy type of thing. So hydrogen bonding when it's enforced can be really, really strong and so here we have the neat IR spectra of two different isomers of orthohydroxy methyl benzoate

or methyl salicylate and of the meta isomer and the one thing that I thought was kind of cool is the difference between the two OH groups so the OH stretch is at 3190 in the ortho compound and at 3370 in the meta compound.

So these are both hydrogen bonded OHs but in the ortho compound you have enforcement of hydrogen bonding through an intramolecular hydrogen bond and that intramolecular hydrogen bond is really strong

and so you end up being even lower in frequency than the meta compound where you just have intermolecular hydrogen bonding.

All right, well I think that's what I want to say about IR spectroscopy. We'll pick up next time talking about mass spectrometry and we'll spend three lectures. We'll start by talking a little bit about the theory and the instrumentation and then some of the concepts.