Functional Groups

Functional groups are structures within molecules that contribute to the properties of that molecule. In his excellent book, The Same and Not the Same, Roald Hoffmann cites the concept of the functional group as something in chemistry that is not reducible to the physical laws that affect it. Functional groups as concepts seem to be uniquely chemical. I, at least, can't think of anything that's more than superficially analogous. And for organic chemistry, they're hugely important.

C—C bonds and C—H bonds are extremely common in organic molecules. If you remember my post showing how skeletal structures work, you should recall that these bonds are only noted with points and intersections for the former and are left to inference with the latter. Such structures are "skeletal" because they really do show the hydrocarbon skeleton of a molecule. All those C—C and C—H bonds are usually pretty stable. They can contribute to the chemistry of a molecule, but not nearly so prominently as functional groups.

I won't attempt to list every functional group here, because there are lots and lots of them and your puny brain would probably die or something. But I will list some of them. First, I want to introduce a new notation. Actually, I'm not sure if I already introduced it, but I'm too lazy to go back and check and you probably forgot about it anyway. The letter "R" is often used to denote the rest of a molecule apart from a functional group, especially if what remains is a plain old hydrocarbon. This can be convenient for situations when it's only the functional group we care about and drawing the rest of the molecule would be impractical or wouldn't even make sense. If we do this more than once though, and the groups being condensed are not identical, using "R" to denote both would be inappropriate, but "R" for one and "R'" (R prime) for another is fine.

So for an example, I've decided to use 2-butoxyethanol because I've used it to clean graffiti off the walls in the restroom at work.
Of course, while this is the stuff I was using, other molecules that change one or more atoms are easily possible. What if we decided only to focus on part of the molecule? We might do this...
What is "R"? Is it still a chain of four carbons with nine hydrogens? Maybe. Or maybe it's still a straight chain, but one carbon longer than that now. Maybe there's a ring. Maybe it has multiple branches. Maybe there are dragons. No one knows! It's a mystery. Exciting, I'm sure. And that's how "R" works. It's a placeholder. It saves space. "Here be dragons" might work too, but "R" is the standard. I have no idea how it became that way, actually.

Aliphatic hydrocarbons
This term comes from the Greek aleiphas, meaning "fat." Actually, many of the properties fats have come from the long hydrocarbon chains they possess. If an aliphatic hydrocarbon has no π-bonds (that is, no double or triple bonds), it is an alkane. Branches and rings might be present and do affect the properties of alkanes, but they're still called alkanes, although a compound with a ring might be said to be a cycloalkane.

But if π-bonds are present, no matter how few there are or how big the rest of the molecule is, it's not an alkane anymore. A double bond means that the molecule is an alkene. So using what we learned earlier, a molecule that contains this functional group: R₂C=CR₂ (different groups are all "R" here because noting them all with superscripts would be ridiculously clunky) is an alkene. The double bond counts as a functional group. Specifically it is the alkenyl functional group.

Similarly, a triple bond is a functional group. Something with R—C≡C—R' functional group is an alkyne (no matter how many double bonds or single bonds it has). And this is an alkynyl functional group.

Aromatic hydrocarbons
The only hydrocarbons I know of that are not classed as aliphatic are ones containing aromatic rings. You might be anticipating this from the trend with what I've said regarding alkenes and alkynes, but the presence of even a single aromatic ring in an otherwise aliphatic molecule means that the compound is considered aromatic and not aliphatic. Aromaticity is a tricky concept though, and we're not covering it just yet. So for now, the only aromatic ring we'll concern ourselves with is the benzene ring.
It is not a cycloalkene, even thought it might look like one. I mentioned in an earlier post that this type of ring is resonance stabilized. The electrons are evenly distributed around the whole ring. How this happens and whether a particular ring is aromatic or not are subjects for later posts. But this ring, the benzene ring, is aromatic. If it's treated as a functional group, it's the phenyl group. To abbreviate when I'm using condensed structures rather than skeletal structures, I'll probably use "Ph" for "phenyl." So something with this functional group would be R—Ph. Other people sometimes just abbreviate a benzene ring attached to something else as C₆H₅ but I'll try not to do that so as to avoid confusion.

Another well-known aromatic hydrocarbon with its own common name is toluene. It's just like benzene, but where benzene has six hydrogens, one attached to each carbon, toluene replaced on hydrogen with a methyl group, a carbon bonded to three hydrogens. So it's like R—Ph with "R" being "CH₃" except this is a special case and gets its own name. If toluene acts as a functional group itself, with one of the hydrogens on the carbon outside the ring being replaced by a bond to the rest of molecule, and I'll draw a picture just to be sure you're following me...
The functional group is actually not phenyl. It's a benzyl group. If this confuses you, good. I like confusing you. Try to remember that this is a benzyl group and that a benzene ring attached to something is just a pheny group. I know both "benzene" and "benzyl" start with "benz-" and you're really tempted, but don't. Just don't.

Now for some more functional groups...

Alkyl halide (halo group)
R—X (where X is a halogen). It could be an alkyl fluoride, alkyl chloride, alkyl bromide, or alkyl iodide depending on which halogen is attached.

Alcohol (hydroxyl group)
R—OH

Ether (alkoxy group)
R—O—R

Amine (amino group)
R—NH₂ (primary) or R₂NH (secondary) or R₃N (tertiary)

Thiol (mercapto group)
R—SH

Sulfide (alkylthio group)
R—S—R'

Aldehyde (carbonyl group)
R—CHO (for those of you who can't deal with condensed structures very well, the carbon is double-bonded to the oxygen, bonded to the hydrogen, and bonded to the R group, so it could also be R—CH=O)

Ketone (another carbonyl group)
R—CO—R' (or R—C=O—R'). The difference between an aldehyde and a ketone is that in an aldehyde, the carbobyl group is on the end of a chain, but in a ketone, it's in the middle of a chain.

Carboxylic acid (carboxyl group)
R—COOH (like an aldehyde with the hydrogen on the carbonyl carbon replaced by an hydroxyl group, alternatively it's like an alcohol with an oxygen double-bonded to the terminal carbon)

Ester (ester group)
R—COOR' (like a carboxylic acid, but with the hydrogen on the oxygen replaced by a hydrocarbon group).

Amide (carboxamide group)
R—CONH₂ (primary) or R—CONHR (secondary) or R—CONR₂ (tertiary). Not to be confused with amines, which don't have that oxygen double bonded to the carbon that nitrogen is attached to.

What's that? You want more? Fine. Next time, I'll post some more functional groups for you.

The Next Chapter is About Organic Molecules and Functional Groups

I am so excited for this. With the huge hiatus between the "Acid Strength" post and the "Aspirin" post, I hadn't bothered to look at what the next chapter of my text book had in store for this project. Would I skip the chapter? Go over some of it and move on as quickly as possible? Functional groups are one of my favorite things in the whole world and I can hardly wait to write about them. All chemistry is interesting, but functional groups are what truly fascinate me and I hope to study them when I go back to school. But even before then, I'll be doing a little reading up on them while writing posts here. In the meantime, have a little patience for me. I really am back and giving this project my attention, but I want this next post to be my best one so far. I love functional groups.

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.

Acid Strength

Acid strength refers to the potential of the acid to lose a proton. Acids that give up a proton easily are stronger. Acids that don't readily give up a proton are weaker. Here, I'll show you a reaction mechanism for an acid dissolving in water. I know you don't know what those are, but if you paid attention the first time I showed you, then you would. It's not my fault you weren't paying attention. Grow up. So yeah, mechanism...
For a little review, I'll note that the first reactant (A—H) is the acid. Water, the second reactant, is the base. And the products are the conjugate base of the reactant acid and hydronium, the conjugate acid of water, respectively. This reaction is reversible, so the hydronium could protonate the conjugate base of the original acid and leave us with the reactants again. But the equilibrium favors whichever side of the equation has the weaker acid. This might seem intuitive, actually. The acid that more readily gives up a proton (the stronger acid) will do so more and therefore will show up less. But unless we know which acid is weaker, that is, unless we have a way to measure acid strength, the knowledge doesn't really help us. Fortunately, quantifying the acid strength is possible. But this goes a bit beyond the scope of the textbook I'm using, which assumes the student remembers certain things from general chemistry. Besides, it involves math. So I'll just tell you that in organic chemistry, the figure that is typically used is pK_a. That's the opposite of the common logarithm of the acid dissociation constant. The acid dissociation constant is determined by multiplying the equilibrium concentrations of the products and dividing this by the equilibrium concentration of the acid. Easy, right?

So I totally don't remember how those equilibrium concentrations are determined (with instruments, I guess). I just look up the pK_a of the acid in question using a table of pKa values. That way, all I need to know is that the lower the pK_a, the stronger the acid. My textbook states that typical pK_a values for organic acids range from 5 to 50. I'm not sure where those numbers came from, but whatever. Be aware that some organic acids are not typical. Also, in lab, I dealt with inorganic acids all the time. Acids with negative pK_a values are considered "strong" acids. Actually, Wikipedia says that strong acids are those with pK_a values less than -2. Whatever. What it means for an acid to be "strong" is that essentially all of it will lose its protons to water. In other words, the equilibrium completely favors the products and none of the original acid is present in any concentration. So really, the strongest acid that exists in water is hydronium. Anything stronger just protonates the water to form hydronium.

Acid-Base Definitions

Do you know what acids and bases are? Seriously? I know I used to think I did and I totally didn't. I mean, I knew some examples of acids and bases, but the chemistry behind them was completely unknown to me even after I took chemistry in high school and even when I started taking it in college. But finally, in one class I was introduced to three acid-base definitions. I knew there were more, but I had no idea how many. This post will introduce some of them. Really, I've only ever used two of them and those two, which happen to coincide with each other a lot, are the only two that will be important here, but I think the historical definitions are interesting and this is my blog or whatever and so I get to make a post including them.

Lavoisier definition
In case you didn't know who Lavoisier was, he was one of the founders of chemistry and you are not worthy. Through a series of experiments, he determined that oxygen combining with other elements could lead to "acidic" (the word comes from the Greek word for "sharp") properties. He extrapolated from this that oxygen was the element that contributed the acidity. This is where oxygen got its name, which means "acid-generating." There was a small problem with this: Lavoisier was wrong. Considering that he basically invented chemistry, I think he's allowed to be wrong every once in a while.

Liebig definition
You know who Liebig was too, right? Because he was another great chemist. Anyway, this was sort of the first real definition. A Liebig acid is a molecule that contains at least one hydrogen atom that can be replaced by a metal. This was only a definition for acids. Back then, bases were loosely defined as the opposites of acids. Known acids were generally liquids, so a base was whatever compound would react with an acid to neutralize it into a solid salt. Such reactions are known as acid-base reactions and we'll deal with them other posts pretty soon.

Arrhenius definition
This was one of the three definitions I originally learned and it was the main definition everyone used for a long time. But that was in the past. The distant past. Like before I was born, even. Probably before you were born too. Arrhenius acids are molecules that, in water, lose hydrogen atoms, generating hydrogen ions in the solution. Well, we now know that they're actually hydronium ions. Oh, hydronium is important. You should know what it looks like. Here, I'll paint one for you.
If you weren't stupid, you'd remember what a water molecule looks like. Man, I'm not posting water again here just for you. Go back and find the post where I did show it or something. Anyway, this is like water, but with another hydrogen. Oxygen normally only forms two bonds. I guess I never talked about formal charges or whatever, but the oxygen is positively charged now. Really. We'll talk about it later if you want. Hydronium is properly written as H₃O⁺. But you should be aware that a common shorthand is just to just write H⁺.

So those are Arrhenius acids, but there are also Arrhenius bases. They are molecules that, in water, generate hydroxide ions. You want a hydroxide ion? Here you go.
Note that when it has oxygen has one too many bonds, it is positively charged, but when it has one too few bonds, it is negatively charged. Unlike H⁺, hydroxide ions actually can and do exist in water. The shorthand for hydroxide is ⁻OH. That's how I would notate it if I were writing stuff down by hand, but superscripts and subscripts, although necessary for chemical notation, can be annoying to do on Blogger, so I'll probably stick to just typing "hydroxide" most of the time.

Now might be a good time to mention that in aqueous systems (systems with water as the solvent—I'm not going to elaborate on this for now), hydronium and hydroxide act as a sort of currency of acidity and bacisity. When a reaction takes place, the actual atoms from the acid/base aren't the ones that are participating. I'll illustrate this with a classic acid-base reaction...

HCl + NaOH → NaCl + H₂O.

So, hydrochloric acid and sodium hydroxide yield sodium chloride and water. Sodium chloride is the salt in this case. In other reactions, other salts would be formed. A salt and water are the products of Arrhenius acid-base reactions. But keep in mind that water is the solvent in which this whole thing is taking place. HCl is a gas and NaOH is a solid. When we do this reaction in a lab, we're likely to just mix samples of water that have the compounds dissolved in them. The bond between the chlorine and the hydrogen breaks and the hydrogen reacts with a water molecule to form hydronium while the chlorine atom becomes a chloride ion and sits there dissolved in the water. The bond between sodium and oxygen likewise breaks and we're left with hydroxide and a sodium ion, which also sits there dissolved in the water.

The hydronium ion produced by dissolving the acid in water can and will react with another water molecule. But the product of that reaction leaves the original hydronium ion turned into an ordinary water molecule and the water molecule attacked by hydronium as the new hydronium ion. This process occurs repeatedly, but without really changing anything because the number of hydronium ions remains constant. The same is true with the base. If hydroxide reacts with a water molecule, the hydroxide gains a proton and is now a water molecule while the water molecule it reacted with lost a proton and is now a hydroxide ion. The charges are rapidly transferred from one water molecule to the next, but they remain there. This tranfer would continue for a long time, but when we mix the two solutions, in addition to reacting with water, the hydroniums and hydroxides can react with each other. Whenever this reaction takes place, the two ions neutralize each other and we're left with only water (H₃O⁺ + ⁻OH → 2H₂O). This reaction also produces heat, so if you do it at home for some reason, keep that in mind so that you don't die or whatever. If you were wondering, the sodium and chloride ions stay dissolved. You have seen what happens when you add sodium chloride to water, right? Seriously, you'd better have. If not, go do it right now.

So that's the Arrhenius definition. There are several other definitions, some of them relevant to certain fields, but two definitions are by far the most popular and important, so I'll introduce them now.

Brønsted-Lowry definition
This is the definition that seems to be used the most in my textbook and introductory chemistry courses. It is more inclusive than the Arrhenius definition. A Brønsted-Lowry acid is a proton donor. A Brønsted-Lowry base is a proton acceptor. One big difference between this and the older Arrhenius definition is that water is not necessarily present as a solvent. It can be and often is, but it's not necessary. The most important difference to keep in mind might be that not every Brønsted-Lowry base will lose hydroxide. Every Brønsted-Lowry acid must have a hydrogen atom that can be lost in order to give off a proton, but the bases need not look anything like Arrhenius bases. What a Brønsted-Lowry base does need is the ability to form a bond to a proton. That means it needs free valence electrons that can be recruited to form the bond. Lone pairs of electrons work best, but electron pairs in π-bonds also have basic potential.

My next few posts will be dealing with Brønsted-Lowry acids and reactions with them, so this is the definition to pay the most attention to. Hopefully, you'll get a feel for what chemicals act as good acids and bases and how acid-base reactions work. For now, definitely remember that an acid is a proton donor and a base is a proton acceptor.

You might have deduced that once an acid has donated a proton, what's left is capable, under the right circumstances, of accepting a proton and recreating the original acid. Likewise, a base that accepts a proton now has a proton that could be donated to something else, recreating the original base. These are conjugate acid/bases. That is, whenever an acid-base reaction occurs, the acid becomes its conjugate base and the base becomes its conjugate acid. The reaction could potentially reverse. Here's one reaction...

H—A (acid) + :B (base) → :A⁻ (conjugate base) + H—B⁺ (conjugate acid)

Now we'll reverse the reaction...

:A⁻ (conjugate base) + H—B⁺ (conjugate acid) → H—A (acid) + :B (base)

The system can bounce back and forth between these two. So where does that leave us? The rule I learned in general chemistry and used all the time in my organic chemistry classes was that the equilibrium favors the side with the weaker acid. Unhelpful if you don't know which acid is weaker, but we have ways of figuring that out.

Under the Brønsted-Lowry definition, acidity and basicity are relative to what something is reacting with. Water, for example, is an acid when it's reacting with hydroxide (it donates a proton) and a base when reacting with hydronium (it accepts a proton). Also, you might have deduced that the conjugate acid of water is hydronium and the conjugate base of water is hydroxide. If you deduced that, good job. You get a gold star.

Lewis definition
This definition was formulated by Gilbert N. Lewis. You do know who he was, right? A Lewis acid is an electron pair acceptor. A Lewis base is an electron pair donor. This definition is more inclusive than the Brønsted-Lowry definition. A base still needs free valence electrons to donate, but the acid in the reaction no longer needs to provide hydrogen. So while all Lewis bases are also potential Brønsted-Lowry bases, not all Lewis acids are Brønsted-Lowry acids.

I'll say more about Lewis acids later. This material is in the last section of the acids and bases chapter in my textbook and it's time to move on to some specific Brønsted-Lowry acid-base reactions.

Acids and Bases

I finished Chapter 1 of my textbook in March and I should have started Chapter 2 some time ago. I don't understand why Chapter 2 is the one on acids and bases. It looks like it's sandwiched between Chapter 1 and Chapter 3 because this material needed to go somewhere early in the book and there was nowhere else for it. But maybe I'm wrong. I'll cover this material, just because I'm not sure.

My textbook is only covering two definitions of acids and bases. But there are others. Looking at Wikipedia, there are some I've never even heard of. I think understanding these would be useful, and it's something that's new to me. This means a bit of a delay in this blog getting actual content because I'll have to do some research online.

Mixed Melting Point

I am back from hiatus and I resolve to post here more often. If you follow my Livejournal (you know, the only Livejournal that belongs to Stephen Bahl) you already know that I have been slacking off on internet-related things lately because I've been reading actual books printed on paper. I know, it's weird. So that's the main reason I haven't updated here in well over a month.

I notice that my last entry ends with this ...

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.

What? Why did I say I'd explain that in my next post? I think I knew that I would go on hiatus with this project and put that there as a way of punishing myself or something. Well, I suppose I could explain this. Using pictures would help, but I'm going to try to do it without pictures because I like words so much and also because I'm lazy. Here we go...

The melting of a crystal involves the excitation of the particles forming the crystalline structure. In the case of organic compounds, the particles are organic molecules and the crystal is a specifically shaped arrangement of these molecules based on the intermolecular forces at work in the molecules. Now, the reason I say that I am punishing myself by explaining this now is that I haven't yet done a post on intermolecular forces yet. It's not a particularly difficult topic, but rather than going into the details, I'll just simplify and state that intermolecular forces cause molecules to interface with one another. This has a direct effect on the phases of matter (solid, liquid, gas, and the others). A crystal is solid because the intermolecular forces are strong enough to make the molecules stick together in the crystalline pattern. Adding heat excites these molecules and the intermolecular forces are no longer strong enough to hold them together, so the molecules slide against each other and bump around making a big mess or something.

The change in phase from solid crystal to liquid mess is, as you hopefully already know, referred to by the scientific term "melting." The precise temperature at which a pure crystal of a given compound melts is referred to as the compound's melting point. Different compounds, because they have different intermolecular forces based on their structures, have different melting points. So let's say we have Compound A with a melting point of 160°C (that's 320°F for those of you who still don't know the Celsius scale).

We have a sample that we think is Compound A. We can melt it and see that it melts at 160°C just like Compound A should. But now let's say that this sample is actually Compound B. Just by luck, Compound B, although made of completely different molecules than Compound A, happens to also have a melting point around 160°C. If we do a proper melting point analysis, we won't be fooled. Here's why.

We take some of Compound A and some of Compound B, mix them together, stir the powder up a bit and put some of our A/B mixture into a melting point sample tube (a thin tube made of glass, like I desribed in the last post), then use the melting point apparatus. As noted, this is a mixture. We don't have a pure, uniform crystal of one molecule, but pockets of two different molecules jammed up against each other. They interfere with each other's intermolecular forces and rather than a sharp melting point at 160°C, we'll probably see regions of softening and melting appear over a broad range of temperatures well below 160°C (because some areas will have a more even mixture meaning little crystalline structure while other regions might be isolated pockets of nearly pure A or B).

This technique is called a mixed melting point. Makes sense, right?