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.

Electronegativity

That last post was pretty long, huh? I'll bet it was too long for you. Well, you're in luck. This post is going to be a short one because I'm tired and electronegativity seems pretty easy to explain. It should be easy to explain, even to someone like you. That's why I'm skipping some stuff in this book that looks like it's about molecular orbital theory. Do you think you can handle molecular orbital theory? Yeah, maybe you can. I don't know. But for now, we're not going to worry about it.

Electronegativity is a property atoms have. Atoms that are highly electronegative have a high affinity for electrons. Atoms that are not very electronegative have a low affinity for electrons. Is that simple enough for you? Here. I'll personify the atoms for you, just like every high school science teacher on the planet has ever done. Atoms that are highly electronegative want electrons. They will take them. Atoms that are not very electronegative are more likely to lose their electrons to those mean, electronegative bullies. Now I feel all dirty.

Anyway, electronegativity is inversely proportional to atomic radius. So smaller atoms tend to be more electronegative than bigger ones. There are various explanations for this, but I took general chemistry a long time ago. I usually look at it in terms of "shielding" although technically, there's more to it. For an example, let's compare oxygen to sulfur. They're in the same group and have similar chemical properties, after all. Now, oxygen has eight protons, but sulfur has sixteen. And if you were good and read the second post I made here, you know that protons reside in the nucleus of the atom. So all those protons in sulfur are presumably twice as attractive to any electrons as the only half as many protons in oxygen. If that were the whole story, we'd expect sulfur to be more attractive to electrons. But sulfur also has twice as many electrons sitting in its orbitals. And they take up more space. Hence the mention of atomic radius. The sulfur atoms doesn't just have more subatomic particles. It is spatially bigger. And that large, negatively charged space has a repelling effect on other electrons. It's not strong enough to override the attractive effect the nucleus has on electrons. It doesn't shield the nucleus that much. But the effect is enough to make sulfur's electronegativity lower than oxygen's.

Of course, not all elements are in the same group as oxygen. And that matters because the number of valence electrons is also very important when it comes to electronegativity. One very simple way to look at electronegativity is that it increases as one moves to the right on the periodic table (ignoring the very last column for reasons that are obvious if you take general chemistry) and decreases as you move down the periodic table. The most electronegative element of all is fluorine. Second is oxygen, followed closely by chlorine. Fourth is nitrogen followed closely by bromine, with iodine being sixth. Whenever one of these elements forms a covalent bond with some other element I didn't just name, the electronegative one will have more of the electron density from the covalent bond.

By the way, carbon is more electronegative than hydrogen, but only by a little bit. It's not usually enough to concern us. But when oxygen or nitrogen (or any of the halogens) is in a molecule, electronegativity becomes important.

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."