Hydrogen Bonding

Hydrogen bonding is a special (and awesome) intermolecular force that can happen when hydrogen is attached to nitrogen, oxygen, or fluorine. All three of those are pretty small atoms and are also highly electronegative. They pull so much electron density away from the tiny hydrogen that the nucleus (a single proton) of the hydrogen is highly exposed and attracted to negatively charged things, letting it sort of stick to negatively charged bits of other molecules.

Fluorine is a halogen. I wrote about this, but it was so long ago that you probably forgot all about it. As a halogen, fluorine can only have a bond to one other thing. If that thing is hydrogen, then we have lots of hydrogen bonding fun, but there's only one compound for which this is possible, and that's HF (hydrogen fluoride), which isn't organic, so we won't be paying much attention to it right now.

Nitrogen and oxygen, unlike fluorine, can both be attached to hydrogen and have at least one bond to spare. So there are lots and lots of compounds that exhibit hydrogen bonding. Just look for hydrogen attached to nitrogen or oxygen. Hydrogen bonds are stronger than dipole-dipole forces and other intermolecular forces. In fact, they're partial covalent bonds, although still not nearly as strong as regular covalent bonds.

The classic picture for showing hydrogen bonds is with water, so I might as well just use one of those rather than trying to make my own picture. It's much easier and looks way better. Here's one...
Pretty much any property of water has something to do with hydrogen bonding. And it's important in lots of other compounds too. The structures of proteins and nucleic acids use plenty of hydrogen bonds. Hydrogen bonding is one of the reasons hydrogen is my favorite element. But just remember, it's stronger than other intermolecular forces.

Dipole-Dipole Forces

Do you remember how electronegativity works? I posted about it, so you should. Here, just for you, I'll link back to that post. Electronegativity is necessary to understand polarity. And polarity is how these forces operate. An organic molecule with a functional group that contains an electronegative atom such as oxygen or nitrogen will likely have a permanent dipole. The heteroatom pulls negative charge toward itself (because it's electronegative, obviously). As described in the previous post, regions of the molecule can have differential charge. But a permanent dipole is much stronger than the fleeting changes responsible for London dispersion. Because of this, molecules that have dipole-dipole interaction experience stronger intermolecular forces than ones that have only London dispersion. Compounds with this property are said to be polar and ones that do not are non-polar. Consider these examples...

Acetone
Condensed structure: OC(CH₃)₂
It has a permanent dipole that looks like this.
Those Greek letters represent partial charge. The electronegative oxygen pulls electron density toward itself, so it is the center of negative charge. The region opposite it, lying near the central carbon, is the most positively charged region of the molecule.

Carbon dioxide
Condensed structure: CO₂
This time, when we draw the arrow through one oxygen, the other oxygen cancels it out. Despite having a highly electronegative element, carbon dioxide is non-polar.

Whether a solvent is polar or non-polar tells chemists a lot about its potential uses, and some reactions need one type of solvent or the other. Acetone is a well-known polar solvent, but water is the best known in this category. Many organic compounds such as n-hexane are commonly used as non-polar solvents. That's all for these at the moment. Just remember that in order for dipole-dipole forces to occur, the compound(s) must have permanent dipoles, and that dipole-dipole forces are usually much stronger than London dispersion. Of course, a molecule can have both. But the stronger forces are considered to override the weaker ones for all practical purposes that I've encountered.

London Dispersion Forces

My next few posts will discuss intermolecular forces. The textbook put this information in the same chapter as functional groups, because functional groups play such a huge role in the type and strength of intermolecular forces a compound has. I don't know if this is the best approach, but I'm too lazy to figure out a different one and whatever, this works.

Hopefully the word "intermolecular" tips you off to the fact that these are forces that occur between molecules. If not, what is wrong with you. Intermolecular forces are much weaker than normal covalent or ionic bonds. If they weren't, they wouldn't be intermolecular because they'd be binding atoms as tightly as the bonds within the molecules and the molecules wouldn't really have distinct identities at all.

London dispersion forces are the weakest and most ubiquitous ones we'll look at. In this textbook and in the college chemistry classes I took, they're more often referred to as van der Waals forces. But from what I can tell, that term is actually more inclusive, so I'll be calling them "London dispersion forces" or "dispersion forces" or maybe "London forces" until I have some reason not to. Do note that it seems to be common to refer to these as "van der Waals" forces/interactions. They're also known as induced dipole–dipole forces.

London dispersion is so ubiquitous in organic chemistry in part because it's the basis for intermolecular attraction in hydrocarbons. The other forces don't come into play unless there are functional groups. It would be nice to have the picture in my textbook to convey what I'm talking about, because it really does seem better than what I'm finding on the web, but this ball & stick picture that uses hydrogen molecules should serve well enough...
Because of electromagnetic repulsion, positive and negative charge gather in regions of the molecule. The whole molecule is neutral in charge, but some parts of it are more positive and others are more negative. I like the methane example used in my textbook better, because it shows more randomness, but the principle is the same: the negative portion of one molecule is attracted to the positive portion of another molecule. And this interaction, magnified over massive numbers of molecules, gives the whole thing some cohesion. And when I say some, in the case of hydrogen at normal temperatures and pressures, it's not enough. That's why pure hydrogen is usually a gas. The same is true with methane. If it gets cold enough, as happens in that atmosphere of Titan, one of Saturn's moons, methane is liquid.

Not all hydrocarbons are gases at room temperature. The bigger they are, the higher the temperature has to be in order to excite the movements of the molecules enough to overcome London dispersion forces and liberate the molecules from each other. But it's not all because the molecules are simply bigger. They're also longer. And that means more surface area. My textbook points out that n-pentane has stronger dispersion forces than neopentane. Condensed structures should be enough to show why...

n-pentane: CH₃CH₂CH₂CH₂CH₃
neopentane: C(CH₃)₄

The long, straight chain of n-pentane has more surface area than the bunched up neopentane. They have the exact same quantities of each atom, but this difference has significant effects on the physical properties of these molecules.

Finally, dispersion is also affected by the polarizability of atoms. In a large atom, the electrons are further from the nucleus and are more subject to disruption and fluctuation in electromagnetic forces than in a small atom, where the electrons are more tightly held.

Oh, and a fun trivia fact that everyone loves is that geckos hold onto surfaces using these forces (the bristles on their toes are small enough to do this and well distributed enough that enough surface area is there for the forces to be strong enough to hold the lizard's whole body even when walking upside-down). Seriously. Pretty cool, huh?

Visualizing Functional Groups

I told you that I would post more functional groups and I meant it. But I also want to make the ones I've already introduced clear. And condensed structures can confuse people. I heard you're easily confused. So we'll spend a bit more time getting acquainted with these. I think I covered the hydrocarbons well enough, so we'll focus on functional groups that contain heteroatoms.

More about alcohols
The nature of the carbon that the hydroxyl group is attached to determines the type of alcohol here. First, there's a primary alcohol...
Note that the carbon attached to oxygen is attached to only one other carbon (in the R group). In a secondary alcohol, this carbon is attached to two other carbon atoms...
And as you might have guessed, in a tertiary alcohol, that carbon is bonded to three other carbons...
If you were thinking that a quaternary alcohol would be one in which that carbon is attached to four other carbons, you sure are dumb. Carbon is tetravalent. You remember that, don't you? It can't form five bonds. There is no such thing as a quaternary alcohol.

If you were wondering, whether an alcohol is primary, secondary, or tertiary has important implications for its chemical properties, hence the distinction. This is also the case with amines, as I already alluded to, but in that case, it's how many carbons the nitrogen is attached to that determine which type of amine the molecule is. So there's some cool new information for you. But now for some clarity on material I already covered in my last post.

Visualizing aldehydes & ketones
Here's the Lewis structure of an aldehyde...



And here, for contrast, is a ketone...
Notice the big difference: with an aldehyde, the oxygen is at the end of a chain and with a ketone, the oxygen is attached to a carbon that is somewhere in the middle of a chain. These structures both have a "carbonyl" group and their chemical properties are often similar, but they can be different in important ways and this distinction is certainly worth remembering.

Carboxylic acids and friends
Carboxylic acids get several other classes of compounds grouped with them as "derivatives of carboxylic acids" quite literally because carboxylic acids can be used to make these other compounds. I won't cover all of them because there's a whole chapter on this stuff and it's way later in my textbook. But because you're slow, I worry about your ability to even deduce the general appearance of these groups from a condensed structure. So here's a carboxylic acid...
Like the aldehydes and ketones, there's a carbon double-bonded to an oxygen and single-bonded to an R-group. But the fourth bond isn't to hydrogen or another carbon. It's to oxygen, which itself is attached to hydrogen. Remember acidity? You know, that thing the last chapter was all about and such. And maybe you even remember that in my "Aspirin" post I said, of the carboxylic acid, "This arrangement of atoms makes it easy for a certain reaction to occur. That reaction is a Brønsted-Lowry acid-base reaction." Really, it's not familiar. Whatever. That proton can totally come off.

Since I like functional groups so much, here are some more in condensed structure...

Acyl halide
R—COX (like a carboxylic acid, but with the second oxygen replaced by a halogen)

Imine (imino group)
R=N—R' (these come in multiple varieties and I haven't really studied them yet)

Peroxide (peroxy group)
R—O—O—R' (the oxygens are actually attached to one another)

Nitrile (cyano group)
R—C≡N

Enough. We will now cover new functional groups as they come up.

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.