Strength of Intermolecular Forces

This chapter in the textbook is quite long, but not all of it is well-suited to posts like these. A lot of this reviews concepts from general chemistry and has lots of pretty pictures and I don't want to spend too much time on things like melting point and solubility and soap. The soap thing is something I originally learned in high school and got to see repeated in two general chemistry classes in college and organic chemistry too. I might do a post on it, but for me, it's gotten kind of old. There is some really great material here. I especially like the explanations of biomolecules, but perhaps that's best reserved for later.

In short, I do want to write at least one more post on the odds and ends in the third chapter of my textbook. They will come soon if at all, because I am long overdue on starting the fourth chapter. Before I do either of those things, let's wrap up intermolecular forces.

The strength of intermolecular forces is, in ascending order...

• London Dispersion: caused by fluctuations in charge density across the surfaces of molecules.
• Dipole-Dipole: caused by permanent dipoles.
• Hydrogen Bonding: Caused by extreme loss of electron density on hydrogen when bonded to oxygen or nitrogen (or fluorine, technically).

Note that ion-ion forces, which hold ions together in ionic compounds, could be compared to these forces, although ions are technically not considered molecules. Ion-ion forces are much stronger than any of these intermolecular forces.

Intermolecular forces have the following effects on physical properties...

• Stronger intermolecular forces increase boiling point.
• Stronger intermolecular forces increase melting point.

There's more, but I am skipping it because that's the important stuff and you'd forget the rest anyway. If something I omitted here becomes important later, I'll just blame the problem on you. It's either that or explain the thing when the issue comes up.

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?