Earlier today I said that I was going to do a post on isomers tomorrow. Well, I'm so anxious to talk about isomers that I am starting it early. Only half an hour left until tomorrow anyway as I'm typing this sentence, so maybe I won't finish it until it's tomorrow. We'll see.
Constitutional isomers are compounds that have the same atoms, but arranged in a different structure. This is completely different from resonance structures, because both isomers are real molecules and can have very different properties. So don't get the two mixed up. Remember, resonance structures have double-headed arrows between them. Isomers don't. Because I'm lazy, I'll just use the first example I find in my book... Both molecules have the same atoms. They each have three carbons, one oxygen, and six hydrogens. But in molecule A, the double bond is between two of the carbons and the oxygen is bonded to a hydrogen. In molecule B, the double bond is between the oxygen and the carbon it's attached to, while the other carbons each have three hydrogens.
Remember how it's bonding that really makes molecules what they are? These compounds have different bonding structures, so I would expect them to have very different chemical and physical properties.
Addendum: People should know what common or important molecules look like. I just realized, as I was about to close my textbook, that "molecule B" is, in fact, acetone, a chemical you might be familiar with. Acetone is an important solvent. I used it in the laboratory all the time in my chemistry classes. It's also in paints and stuff. And it's used to make acrylic glass. As for the other molecule (A), its name is 2-propenol and I'm pretty sure it's just an enol form of acetone, so basically it is highly unstable and will typically turn into acetone by itself. But that's a topic for another day.
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.
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.
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.
Both molecules have the same atoms. They each have three carbons, one oxygen, and six hydrogens. But in molecule A, the double bond is between two of the carbons and the oxygen is bonded to a hydrogen. In molecule B, the double bond is between the oxygen and the carbon it's attached to, while the other carbons each have three hydrogens.
