Diastereomeric Alkenes

Yes, it's a new post. I know. I hardly posted at all this year with school, then finished that in June, made two posts, and disappeared. No posts in July or August. It was not my intention to put this project on hiatus. I am sorry I failed you. I don't even know what kept me away for most of this time, but recently, there has been one thing: I am planning a trip. It's a pretty big one, actually. I'll be going to Germany. Since this is going to happen fairly soon and since my German is not so good, working on becoming as fluent in the language as I can be takes priority over writing new posts here. I guess this means I have to go on an actual, planned hiatus. Or, if you prefer, this means I'll have to extend the current hiatus, albeit interrupted by this solitary post. Whatever.

My primary reason for choosing German as a language to learn in the first place was its connection to chemistry. I haven't written about that here, so I'll explain. Germany as a region has long been a powerful contributor to the body of scientific knowledge. This was especially true in the nineteenth century, a formative period for many branches of chemistry. Because of this, in the twentieth century, in American colleges that required students to study a foreign language, chemistry majors traditionally chose German and were often encouraged to do so because it would allow them to access chemical literature that had only ever been published in German. The utility behind this is probably all but vanished these days, as any text that's of practical use to a chemist is probably available in English. But I enjoy history and the prospect of some day being able to read old chemistry texts in German has some sort of allure for me.

The German classes I took this year in school didn't do much in the way of making me comfortable with the language (the instructor mostly just played videocassettes from the 1990's), but they did make me really want to visit Germany. And so, here we are, with me abandoning you yet again. I'm really quite sorry about that.

Rest assured, this blog will be back with a vengeance. After the chapter I've been covering, there's material that I really like and I am excited about presenting it to you. Just seeing this material in the book makes me brainstorm different ways to cover it. There's some really cool chemistry to come once I finish this chapter, which I wish I'd already done over the summer instead of putting it off. So yeah, I should be back home and caught up on everything else by, let's say, October 29th or so. Mark your calendars for late October. Back with a vengeance. Really. More frequent updates. Better updates. Awesome chemistry. But not yet. You'll have to wait for my return. It will totally be worth the wait, though.

In honor of my trip, I'm going to skip ahead a bit to something with an obvious connection to organic chemistry's German roots. We'll return the to the material we were on shortly after I get back. And really, now might be as good a time as any to talk about this. By now, I'm sure you have a decent grasp of one type of stereoisomerism. But as you might have guessed, having a chiral center isn't the only way for stereoisomerism to occur. There are other ways that have nothing to do with a carbon atom bonded to four different groups. In fact, I initially wanted to write a post introducing all of the different ways for this to happen that I knew of, but I couldn't find a way to make it work. I did, once upon a time, say this, though...

Constitutional isomers often have dramatically different chemical properties. Their physical properties differ too. They might have different functional groups. In contrast, stereoisomers don't exhibit such bold differences. Two compounds that are stereoisomers of each other not only have the same atoms, but the atoms are connected in the same way. Their properties are almost identical. But the spatial positions of the atoms are different.

And a chiral center isn't the only way for that to happen. So, here's one of the other types of stereoisomerism, and it's a lot easier to demonstrate graphically in two dimensions than the type you already know about. There's just one thing I've probably never mentioned here that you need to keep in mind: unlike single bonds, there is no rotation along a double bond. Behold...
Can you spot the difference? Don't get too excited about it: everyone else notices it too. Same atoms. The atoms are connected in the same way, with an A and a B connected to a carbon that is double-bonded to another carbon also connected to an A and a B. But the spatial arrangement is different, no matter how we oriented these in three dimensions (try it if you want). We have a notation system (cis/trans) that makes this pretty easy, as seen in the drawing I just made. The "cis" version of this molecule has both A's on one side of the double bond and both B's on the other side of it. The "trans" version has an A and B on each side. They're diastereomers, which you recall from my last post means that they are stereoisomers, but not mirror images of each other. Easy, right?

But what if instead of just having two different kinds of groups, we have three or four? This notation system doesn't have a way to deal with those situations! For that, we need the E,Z system as my textbook calls it or Z-E Isomerie as the German Wikipedia calls it. And it's at this point that I realize I've gotten ahead of myself. In order to explain this, I need to use information that would come from a post I haven't written yet, probably the post I was supposed to have written if I didn't skip ahead to this section because I remembered it having something German in it. So despite my failure here, this still seems like a good post to end on before my trip.

If I had class, I'd rewrite this post and just make the whole thing be about cis/trans notation and save E/Z for later. Since I don't, I'll find some way to squeeze the German connection into this post. E and Z are really just more comprehensive versions of cis and trans, which themselves come from Latin instead of German: "cis" means "on the same side" or something like that and "trans" means "across" probably. In the E/Z system, each group is assigned a priority, but I haven't yet written about the rules for determining priority. They're the same ones that are used in the notation system for chiral centers, which is yet another Latin-based system. But since this simpler system already has the Latin words, for the comprehensive system, we use the German words. Crazy, I know. Eventually, I hope to show some examples of names of molecules with "E" or "Z" in them. "Z" stands for zusammen (together) and "E" stands for the word entgegen (against). I leave it to you to figure out which one corresponds to "cis" and which one corresponds to "trans."

A Note on Complexity and Isomerism

My textbook has a table with information that I did not include in my last post, but that may improve understanding of isomerism. In case it is not obvious, the number of isomers grows with the size of a molecule. In my last post, I showed the two isomers of butane. Larger alkanes have even more, because with more atoms, there are more ways to rearrange them. Small alkanes are easy to understand in this regard. A hydrocarbon with one carbon has no isomerism. The same is true for two or three carbons. When we get to four, as already demonstrated, there are two possibilities: a straight chain and one with a branch. Five carbons means three isomers. With seven carbons, we get nine isomers, which is still manageable, but then add a single carbon and there are eighteen isomers. The table ends with icosane (C₂₀H₄₂), which has 366,319 constitutional isomers.

And that is just acyclic alkanes. There are so many other things to consider, that the complexity is staggering. And that is why we have a systematic method of naming molecules. Anything else would get pretty impractical.

Notation of Molecular Structure

I'm skipping some stuff in this textbook that's about molecule shape, bond angle and other important stuff. I'll get back to it later. Remind me to do that. Right now, I get to go over drawing organic molecules. Awesome. No really, this will be good to do because once this is out of the way, I can assume that you understand how these structures work and move on to whatever I want to. So it's very important that you understand this material. It will let you understand the notation I'll be using from now on. Fortunately, it's easier than you might think: I'm only introducing two structures here.

Condensed Structures:

Condensed structures are what they sound like. They take molecules like the ones I've been showing in previous posts and convey all of the structural information in a short line of text. They're not really practical in certain cases, but for most small molecules, they're easy to write and understand. I'll be using them when I can because I won't need MS Paint or anything like that. Let's start with some rules...

• All of the atoms are drawn in, but single bonds usually aren't.
• Atoms are drawn in next to atoms they are bonded to.
• Parentheses enclose groups of atoms that are all bonded to the same atom.
• Pairs of electrons are omitted.
• Read the structure from left to right and remember that every carbon must be tetravalent.

Don't get it? Don't worry. Here are some examples.

First example: n-butane
Molecular formula: C₄H₁₀
Lewis structure:









Condensed structure: CH₃CH₂CH₂CH₃

Notice the molecular formula, while telling us what atoms are in the molecule, doesn't tell us where they are. The Lewis structure does that. So does the condensed structure. We don't actually get the bonds drawn for us, though. Starting from the left, we have a carbon and three hydrogens. The assumption is that the hydrogens are bonded to the nearest atom, and since each hydrogen can only bond to one thing, that means that the first carbon is bonded to all three of them, leaving one bond open. Next there's another carbon, which must then be attached to that first one, meaning that the first carbon is tetravalent and the second one must still have three more bonds. Two hydrogens are written in after it, so they're both attached to that second carbon, leaving it with one more bond, and so on. With practice, it becomes quite easy to convert between regular Lewis structures and condensed structures, but you probably won't practice, so I'm not sure how easily this will come.

One note on this condensed structure is that it could be further condensed. When the same sequence is repeated, it's acceptable to enclose it in parentheses and use a subscript to note the number of repetitions. So an alternate condensed structure for this molecule would be CH₃(CH₂)₂CH_3. I don't want to be confusing here, but it's the first example the textbook used. Oh well, it does provide the lesson that there isn't necessarily just one condensed structure for a molecule.

Second example: isobutane
Molecular formula: C₄H₁₀
Lewis structure:
The hydrogen atoms are crowding into each other here, which is why I had to cut off two of the bond lengths. In reality all of the hydrogen-carbon bonds would be identical except for the one on the hydrogen attached to the middle carbon. But that's a topic for a more advanced post. I should have just made the carbon-carbon bonds dashes and the hydrogen-carbon bonds hyphens in order to avoid the crowding, but I already made this stupid picture and I'm not recreating it now, so you're stuck with it.

Condensed structure: CH(CH₃)₃

The molecular formula is the same as the previous example. But the condensed structure is completely different. It shows that the first carbon has a hydrogen attached to it, which takes up only one bond, so it has three left. Then there are three CH₃ groups (that's what the parentheses are for) all attached to the same atom, which must be that first carbon. So in our two examples so far, we've used parentheses in two different ways. In the first example, they enclosed a chain of the same repeated group and in the second example, they enclosed identical groups all attached to the same single atom.

Third example: 2-butene Molecular formula: C₄H₈
Lewis structure:









Condensed structure: CH₃CH=CHCH₃

We still depict multiple bonds, which is why I was overjoyed when I realized that I could use the "≡" symbol, already having "=" to represent double bonds. Everything else in this example should be familiar from the two previous examples. The hydrogens are listed directly after the carbons they're bonded to and each carbon is listed after the carbon it's bonded to.

Fourth example: methyl acetate
Molecular formula: C₃H₆O₂
Lewis structure:
Don't worry about the name here. Nomenclature of molecules will come later. Now, if you stayed awake for the post on Lewis structures, you should understand the structure of this molecule.






Condensed structure: CH₃CO₂CH₃

Are you getting the hang of this yet, or will you be a failure forever? We start on the left with a carbon that has three hydrogens attached, then attach the next carbon to it, which apparently has two oxygens attached to it. But how do we know if they're attached by single or double bonds? Well, like I already said, each carbon must be tetravalent. The second carbon is already attached to the first carbon, so that's one. It can't have two double-bonded oxygens attached to it, because that would be five bonds to a carbon atom. It could have a single bond to each oxygen and a bond to the next carbon, but that would leave both oxygen atoms with only one bond and oxygen forms two bonds (uh, unless it's a free radical or an ion or something, but this isn't, so shut up). The only possiblity left is that the second carbon is double-bonded to one oxygen and single bonded to the other, leaving that oxygen with one more bond left, which goes to the next carbon. And finally that last carbon is attached to three hydrogens. Easy.

That's how the textbook does it, but it's not how my professor did it. He would draw the condensed structure like this: CH₃COOCH_3. This is fine, but I had to be careful, because I kept thinking that this version of a condensed structure was depicting a peroxide, in which the oxgygen atoms actually are attached to each other. So don't make that same mistake. And remember to make sure that the atoms have the right amounts of bonds. If that structure really were depicting a peroxide, it would mean that the second carbon is only forming three bonds.

Skeletal Structures:

I like skeletal structures. They're a lot easier to draw than condensed structures, but they're pretty much impossible to type, as far as I know, so I'll have to use MS Paint or something. Many compounds have rings. Using skeletal structures makes depicting rings easy. They're also good for large, branching compounds and the like. They can also be used for smaller compounds down to ones with only two carbon atoms. Again, we'll go over some rules...

• Any time there is a corner where two line segments meet or a point where a line segment just ends, that represents a carbon atom.
• Hydrogens attached to carbons are not depicted. You're supposed to be able to figure out how many hydrogens are attached to each carbon all by yourself. You can do that right? I mean, what are you, a child?
• Atoms that are not carbon or hydrogen (also known as heteroatoms) are drawn in and any hydrogens attached to them are drawn in too.
• Triple bonds mess with the first rule a little bit because they force the atoms involved to form a straight line, so there won't be any corners, but know that if no heteroatoms are shown there, the things forming the triple bond must be carbons.
• Skeletal structures are awesome.
  

Since you've already mastered condensed structures, I'm not going to waste my precious time drawing the Lewis structure for this first example. Figure it out yourself.

First example: n-hexane
Molecular formula: C₆H₁₄
Condensed structure: CH₃CH₂CH₂CH₂CH₂CH₃
Skeletal structure:
Yeah, I was too lazy to try to make it on MS Paint, so I spent more time than it would take me to just draw that damn thing looking around the web for an image I could use and trying to get it the right size. I still managed to make it pretty small and puny, but you get the idea. In the future, I'll either get this figured out or just go all sloppy and draw my skeletal structures in MS Paint.


Both ends are carbons and each corner is a carbon atom. If you can count, you'll realize that's six carbon atoms all in a straight chain. And if you're not stupid, by now you realize that the carbons in the middle of the chain still need two bonds each, so they'll all be bonded to a pair of hydrogens, while the carbons on the ends of the chain will need three hydrogens. And look, that perfectly matches the condensed structure. And it has the same numbers of both types of atoms as the molecular formula. It's almost as though this is science or something.

Second example: cyclohexane
Molecular formula: C₆H₁₂
Condensed structure: Fool, you cannot draw a condensed structure for cyclohexane (it has a ring).
Fine then, Lewis structure:
Yeah, I know. It looks like crap. My awful drawing skills are preserved for all the world to see. Moving on...







That sure is ugly, so let's see the skeletal structure:









Yeah, that's right. It's a hexagon. I'm pretty sure even you are smart enough to know that a hexagon has six corners. Each one represents a carbon. And since each carbon is bonded to two other carbons, that means there must be two hydrogens on each carbon.

That's enough about skeletal structures for now. They can get a lot more complicated, but I'll try to ease you into it, rather than saying, "Look, it's capsaicin."

Stephen Bahl Has Triple Bonds

So I totally realize that I can use a symbol for a triple bond. Yep. It makes things easier. Like here's ethyne, also known as acetylene...

H—C≡C—H

I chose that molecule because I didn't even need MS Paint to do it. Awesome. I'm so excited to have that symbol at my disposal...

≡≡≡≡≡≡≡≡≡≡≡≡≡≡≡≡≡≡≡≡≡≡≡≡≡≡≡≡≡≡≡≡≡≡≡≡≡≡≡≡≡≡≡≡≡≡≡≡

Awesome. The symbol is actually available because of its meaning in logic and mathematics, not because of anything to do with chemistry. But I'll take what I can get.

Resonance Stabilization

Once again, it's been a while. I had so much fun with the Lewis structures post that I wanted to make a new post the very next day, but other things kept getting in the way. Well, I'm not putting it off anymore. My introduces resonance after introducing Lewis structures, so I suppose I'll follow its lead...

Many molecules, both organic and inorganic, are resonance stabilized. Such molecules cannot be represented by a single Lewis structure. For that, we can use resonance structures—if we want to. I don't really like resonance structures myself. But you know, whatever works. Before I go any further with resonance structures, I must explain why they're used at all. They exist because of the delocalization of electrons. Delocalized electrons aren't sitting on any one atom and they're not locked into any one bond. That's why we use the word "delocalized." It's like they're spread out over multiple atoms as one big forcefield of negative awesomeness. I don't know. Let's just move on to the example.

The example my book uses is this anion. Don't worry about its name because I don't know its name either. It's the conjugate base of formamide and it's an anion, but beyond that, I have no idea. Here's are the resonance structures...
Yes, the circled "-" sign is a negative charge. That was obvious, right? That double-headed arrow indicates that these are resonance structures. The important thing to realize is that these structures aren't two molecules. They're representing the same molecule in two different ways. The electrical charge is in two different places, but all of the atoms are in the same places. The first structure makes the oxygen the center of negative charge. The second structure makes the nitrogen the center of negative charge. In actuality, the negative charge is distributed between those two (and in this case, it will be slightly more centered on the oxygen because oxygen is more electronegative than nitrogen).

Resonance structures are not real. My professor compared it to describing a rhinoceros, to someone who had never seen one, as a cross between a unicorn and a dragon. Yeah, that doesn't make any sense. He realized that after he said it. But he was trying to think of a more everyday example of using two fictitious things to describe one real thing. Resonance structures aren't always in pairs, though. Many molecules have three resonance structures.

So yeah, not real. Resonance structures. The bond between the carbon and the oxygen isn't actually a double bond. And the bond between the carbon and the nitrogen isn't a double bond either. It's more like the bond between carbon and oxygen is a little bit more than a 1.5 bond and the bond between the carbon and the nitrogen is a little bit less than a 1.5 bond. I don't know the actual numbers. It's possible to take a measurement, but I don't have the equipment or the expertise, so shut up. We'll just pretend that it's exactly 1.5 on both, even though I know that it isn't. Actually, why didn't this stupid book just use an example where that was the case? I mean, there are cases like that. This molecule isn't one of them. Anyway, it's more like those bonds are each 1.5 bonds. But Lewis structures don't have a way of representing fractional bonds, so we need a way to represent what's going on here. Some day, I'll fix the notation in chemistry. Until then, we're stuck with resonance structures.

This post was probably kind of confusing. I know I get resonance structures. But maybe you don't. Remind me to help make this post more clear. Right now I need sleep. I'm totally going to write about isomers tomorrow.

Lewis Structures

I am not pleased with the pace at which I've been updating this blog so far. I am still getting the hang of it and I think with more practice I'll do better, but right now it's frustrating. Checking out my textbook again, I'm not sure what to do about the whole "homework" thing. The first chapter has 84 problems. I could do them all. I do have the solutions manual if I get stuck and I generally remember this stuff. That's not the issue for me. I don't want to do these problems. They look too easy. At least from what I've skimmed through, this is really basic stuff. I don't want to plod through it before moving on to the problems I really need to work on. I apparently thought that I would work on problems and write here about the concepts I'm studying. But I'm just not motivated to do 84 problems that will mostly be really easy anyway. Maybe I'll make some sort of split where I work on the problems later in the book and write about the simpler stuff. I don't know if that could work.

Enough about me. The first chapter of this book sure seems to think that Lewis structures are important. So we'll talk about Lewis structures. This would all be so much easier if I could write stuff by hand. The QWERTY keyboard does not, to my knowledge, really do Lewis structures. I know I could easily go over Lewis structures in some sort of classroom setting, but here on the web, I don't have even have a whiteboard. Or do I?
Well, it's a toy I'll need to practice with (practice, practice, practice--it's all about practice).

So with Lewis structures, my book cites three general rules

1. Draw only the valence electrons.
2. Give every second-row element an octet of electrons, if possible.
3. Give each hydrogen two electrons.

Also, a line represents a covalent bond. As aslways, each covalent bond is made up of two electrons. Here's a Lewis structure for methane.


I could make all the bonds the same length, as they are in the real molecule, but I'm really bad at drawing and crap, so you'd better get used to it now. Also note that this is not what the molecule actually looks like. This is just the Lewis structure. The hydrogens repel one another and so the configuration they'll be in is the one where each hydrogen is as far away as possible from each other hydrogen. Last time I checked, the world is not two-dimensional, so the actual molecule ends up with a tetrahedral shape. Each hydrogen is in one corner of the tetrahedron and carbon is in the center. I'll save you some time and tell you that the angle between any two of those bonds is 109.5°. But on paper, it's a whole lot easier to just ignore the whole third dimension thing and draw the Lewis structure. Next, let's try the Lewis structure for water.

The two non-bonding pairs of electrons are represented as dots. They're in pairs because that's how electrons roll. They use the buddy system. Actually, it's because each atomic orbital holds two electrons, but that's not what we're worried about right now. The molecule has a bent shape because non-bonding pairs take up more room than bonding pairs due to greater repulsion force and compress bonding electrons closer together. You did know that, right? Please say you knew that. Anyway, the point of Lewis structures isn't really to represent bond angles. But it is possible to have the atoms line up in a straight line. Here's carbon dioxide.
See? The carbon has no non-bonding electrons, so they can't compress anything and the whole molecule has a linear shape. Both oxygen atoms have non-bonding pairs, but think about it. What would they compress? Yeah, the linear shape keeps those electrons as far from each other as they can possibly get. And that's Lewis structures. How about one more? I'll make it one of your favorite organic compounds.