Aspirin

It's been a while and I need to get a feel for this. I'm also not done with the chapter on acids and bases. So I thought it would be cool to show you aspirin. My book does it, although there's some other stuff before it that I sort of went over, kind of, more or less, already. This will also be an opportunity for me to try to use a program other than MS Paint to render something. We'll see how it works...

Well, that seems to work. I had to make it on this ChemSketch program and then copy the thing into MS Paint in order to upload it here, but whatever. Much easier than drawing everything from scratch.

Now let's make sense of this. There are a few functional groups here, and we'll learn all about functional groups some day. Won't that be exciting? The one we're interested in now is the carboxylic acid group. One of the carbons is attached to an oxygen by a double bond and a second oxygen by a single bond, with that second oxygen itself being attached to hydrogen. This arrangement of atoms makes it easy for a certain reaction to occur. That reaction is a Brønsted-Lowry acid-base reaction. Here's a mechanism...
Okay, those arrows look terrible. I'll have to work on that next time. Anyway, that's the same as the mechanism for the Brønsted-Lowry acid-base reaction I covered in my last post. This reaction can happen in the human body, assuming aspirin gets into the body in the first place. Aspirin is not found in nature. I've sometimes heard it said that aspirin was found in willow bark, but that's incorrect. Willow bark contains salicylic acid. Aspirin is acetylsalicylic acid. The stuff in willow bark would look just like aspirin, but the group attached to that oxygen on top would be replaced with a hydrogen. Aspirin became the analgesic of choice, as opposed to salicylic acid or another derivative of it, because it lacked the irritating side effects of those drugs.

So the reaction can and does occur in the human body. When this happens, the conjugate base, or acetylsalicylate, is formed. The conjugate base is ionic (with a negative charge on the oxygen that gave up hydrogen) and cannot cross cell membranes. Fortunately, this isn't a huge problem because, like other Brønsted-Lowry acid-base reactions, the reaction that turns aspirin into its conjugate base is reversible. My textbook notes that it's the acid that is present in the stomach, and the base that is present in the intestines. This should be pretty intuitive. When conditions are highly acidic, aspirin, which is only a weak acid, even when it does protonate something, will get protonated right away by the acid around it. In basic conditions, there's not much acid around to do the reverse reaction and what is there might be even weaker than aspirin as an acid anyway.

So yeah, aspirin.