Preparation of Salicylic Acid from Methyl Salicylate

I just realized that, although I've skipped a little, I'm done posting material from the first chapter of the textbook. At this rate, it will take me the rest of my damn life to finish the whole book. As much as I'm anxious to move on to the next chapter, right now I'm going to deal with something more important. It's something that I wanted to write about before even starting with material from my textbook, but it occurred to me that you'd need knowledge of notation to follow this. Well, if you actually read the posts about notation of structure, you should now have that knowledge (unless I'm a bad teacher or you're just stupid or something). What I'll be writing about this time is an experiment I did in my first quarter of organic chemistry. I was going to say that it was the very first experiment in the class, but I checked my lab notebook and I was wrong. This was the fourth experiment in the class.

The lab instruction manual has these silly "scenario" sections for each lab. Sometimes they were excessively silly and other times they were actually rather interesting. I'd like to quote this on in its entirety.

The new-age pharmaceutical company Natural Nostrums manufactures drugs from "natural" starting materials. For example, the company manufactures a painkilling drug it advertises as "organic aspirin" starting with methyl salicylate, which occurs naturally in wintergreen oil. Most commercially marketed aspirin is manufactured starting with benzene, a product of petroleum refining. An intermediate in both of these syntheses is salicylic acid.

Natural Nostrums claims that its aspirin, which is supposedly more natural than aspirin made from benzene, has fewer side effects than ordinary aspirin. Critics have accused the company of false and misleading advertising, asserting that salicylic acid made from methyl salicylate is no different than salicylic acid made from benzene, and that the resulting aspirin is no better than any other aspirin.

Les Payne, Director of Operations for the Association for Safe Pharmaceuticals (ASP), is investigating the company. He just shipped your supervisor a sample of salicylic acid manufactured from benzene and a bottle of methyl salicylate that one of his agents obtained from the chemical stockroom at Natural Nostrums. Your assignment is to prepare salicylic acid from methyl salicylate and find out whether or not it differs from salicylic acid made from benzene.

The main purpose of this experiment, in my class, was to introduce laboratory synthesis of one molecule from another (it was our first true synthesis lab—the previous three dealt with identification, extraction, and both purification/identification respectively). But the "scenario" for the experiment touches on what I consider one of the most fundamental concepts in chemistry: the properties of a chemical are caused by the the atoms comprising the chemical and the bonds between them. There's the law of definite proportions, but that's just part of it. Chemistry assumes, because we've tested it countless times, that provided all of the atoms are there in the same proportions and are of the same isotopes, the bonds are in the same places and arranged in the same way, two molecules are not just indiscernible from each other, but identical. Molecules don't "remember" where the atoms making them up used to be. Carbon 12 is carbon 12. Water is water. Salicylic acid is salicylic acid and if I mixed up pure salicylic acid synthesized from benzene with pure salicylic acid synthesized from methyl salicylate, no one could separate the one from the other or tell me which source a randomly chosen molecule came from.

I don't think I've ever seen anyone else articulate just that, even though every chemist knows it to be true. In all my chemistry classes, the closest thing I've seen to anyone pointing out this fundamental fact that really forms the basis for all chemistry is that quote from my lab instruction manual. Maybe people think it's so simple and obvious that it's not worth mentioning. I strongly disagree. Scams like the fictional one described by that scenario really do exist, and people only fall into these traps because they know nothing about chemistry, meanwhile the part of chemistry that makes it clear why these scams are wrong, the same part that's fundamental to all work done in chemistry, isn't considered worth noting.

Then there's homeopathy. Don't get me started on that. Enough ranting though. On with the science! You do want to know how we made salicylic acid from methyl salicylate, right? Of course you do. Well, to begin, here's methyl salicylate.







You might want to just not worry about the ring for now. It's important, but it's not changing into anything else in any reactions I'll be covering for this post. If you are curious, know that the ring has resonance stabilization: the electrons that make up the double bonds in that image are actually evenly distributed across the whole ring. It's called an aromatic ring. That's not because it has a strong aroma. It's for historical reasons. But in this particular case, the molecule actually does have a strong aroma. Methyl salicylate is the chemical that creates the "wintergreen" smell. That's probably a big reason why my professor chose it. When you're doing labs with chemicals that smell bad all the time, ever once in a while it's nice to do a lab with something that smells nice. Anyway, that's methyl salicylate. Here's salicylic acid.








So both of them have that aromatic ring and the hydroxyl group (OH) attached to one carbon with a second group attached to an adjacent carbon. But with methyl salicylate, the group is —COOCH₃ whereas with salicylic acid, the group is COOH. In other words, the difference is that the oxygen the carbon is singly bonded to is attached to a hydrogen or a methyl group, depending on which compound this is. In order to go from one to the other, we need some sort of chemical reaction to change that hydrogen into a methyl group. Chemistry! Oh yeah.

What we do is put melthyl salicylate in a vial, then add an aqueous solution of sodium hydroxide. I wrote a bit about ions a while back, but if you don't remember, ions tend to dissociate from each other in water. The cation (positively charged) is sodium. The anion (negatively charged) is hydroxide, or ⁻OH. At any point, the hydroxide could rip a proton (the nucleus of a hydrogen atom) from a water molecule (this would be an acid-base reaction with water as the acid and hydroxide as the base) and become a water molecule itself, but then the water molecule that it reacted with would be hydroxide, so the total amount of hydroxide ions stays the same.

This mixture is then heated under reflux. What that means in this case is that the mixture is boiling, but there is a glass tube called an air-cooled condenser at the top of the vial. As the vapors from the mixture rise up the tube, the air around them, being cooler than them, absorbs their heat (thermodynamics for the win), and causes them to become liquid again (condensation), at which point they run back down the glass tube into the reaction vial. And all the heat adds kinetic energy to the molecules, which speeds up the rate of reaction. I will now show you the reaction using a mechanism.

Reaction mechanisms show how bonds are formed and broken in reactions using arrows that always start at a pair of electrons and show what those electrons go to. In this case, the electrons from hydroxide form a covalent bond to carbon, breaking the π-bond in the C=O bond, so that the carbon is attached to four things: the ring, the methoxy group (–OCH₃), the hydroxyl group (–OH), and the now negatively charged oxygen that used to be double-bonded, as shown below.But the product of this reaction is an unstable reaction intermediate. It probably exists for only a fraction of a second before the electrons on that oxygen atom form a new bond to the carbon. For reasons that won't be explained right now, the bond that carbon must lose is the one to the methoxy group, so we get this reaction.
Why, look at that. It's almost salicylic acid. Almost. There's just one pesky detail remaining. Being in such basic conditions (lots of hydroxide floating around), the hydroxyl group attached to the ring lost a proton. As long as the conditions are basic, we have no way of putting that proton back on. There's a simple solution to this: douse the whole thing with sulfuric acid. Yes, I'm serious. It's what we did. Well, it wasn't really "dousing." We added more than enough sulfuric acid to neutralize the sodium hydroxide and the disodium salicylate (the name of that last product). At that point, we get salicylic acid. Conveniently for us, salicylic acid is a solid that precipitates out of the solution, so we just filter it out (we use a funnel with filter paper over a vacuum system to pull the liquid through and leave the solid crystals on the filter paper). Finally, we pour cold water onto the filter to wash away any remaining traces of the original solution. Then we just let our crystals dry and we have pure salicylic acid.

We can analyze the purity of our compound by melting point analysis. There are different ways to do this. The method I used was to take three very thin glass tubes (open on one end and rounded at the other) and pressed the open ends of the tubes against crushed powder from crystals, tapping the tubes to shake the samples to the bottoms of them. I put the tubes in a melting point apparatus, basically a heated chamber with a display indicating the temperature and a magnifying glass letting me see the tiny samples in the tubes with ease (I still had to wear my glasses, but shut up). One tube contained my product. One tube (the control) contained salicylic acid from benzene. And one tube contained a mixture of my product and salicylic acid from benzene. For reasons I'll explain in my next post, if my product had not been the same compound as the salicylic acid from benzene, even if the melting points were nearly identical, the tube with the mixture would melt at a lower temperature and over a broad range, as opposed to melting all at once. Since my product was pure, all three samples melted sharply at 160°C.