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.