Intermolecular Forces

Let's talk about intermolecular forces. Intermolecular forces are forces that exist between molecules, not within the molecules. If it's within a molecule, that's actually just called a chemical bond. That would be like one atom attached to another atom. But with intermolecular forces what I'm talking about is one molecule that is loosely associated with another molecule meaning they just like to stick together a little bit.
It turns out that IMFs are what make molecules sticky. I know that that kind of sounds weird that molecules would be sticky, but they actually are. It turns out that if these molecules did not stick together in some way, everything in the universe would be a gas. In order to have solids, in order to have liquids, these molecules have to aggregate. They need to remain somewhat close to each other or they're just going to disperse.
That's what intermolecular forces do. They don't actually change the compound, the compound is affected by bonds, but the way that the state of the matter is affected by intermolecular forces.
I just want to show you guys this really quick diagram. This shows you three different molecules that have about the same molecular weight. Theoretically, if they have the same molecular weight, you would think that overall the type of substance would be similar, but actually, it turns out to be very different.
Here we have a three-carbon atom. I said atom; I meant molecule. Sorry. Here we have a two-carbon molecule with one oxygen. And here I also have a two-carbon molecule with one oxygen. These are very similar looking molecules and they're very similar shapes and yet their boiling points are vastly different.
If you look at propane, propane is like a propane gas. Why do we keep it in a gas tank? Because it's gas. There's no way that you could just put propane in a bowl or in a cup. It would immediately turn into gas.
Then we have dimethyl ether which just by adding one oxygen in that location, we've increased the boiling point a little bit. Remember that ether is one of those reagents that you're going to use a lot in the lab and remember that it's very reactive. And one of the reasons is because it turns into a gas very easily, so it's very volatile.
Then the last one here is ethanol. I've brought up ethanol before. You guys are very familiar with it. That's like vodka. If your vodka was a gas, it probably wouldn't be as easy to consume. But it turns out that ethanol doesn't boil until 173 degrees Celsius.
That's crazy. Look how much different this boiling point is from this one. What's the difference? Why is it such a much higher boiling point? Why does ethanol exist as a liquid in room temperature, whereas propane and dimethyl ether are going to exist as a gas? The reason has to do with intermolecular forces.
Whenever you get questions about boiling points or melting points, that's the way that professors like to ask these questions, it always has to do with the strength of IMFs between molecules. So if a professor ever asks you list the following molecules in order of increasing boiling point or something like that, you know that they're talking about IMFs. 

So there's three main intermolecular forces that we want to know and I'm going to teach you them in order from strongest to weakest. So we're going to start with the strongest, most important intermolecular force and then we're going to end off with the weakest one.
The strongest one is H-bonding, hydrogen bonding. You guys learned about this in gen chem. Maybe you guys can tell me which atoms are the ones that are able to hydrogen bond? Do you remember? It was small, highly electronegative. That was N, O and F. These are the three atoms that if they're attached to an H, they're able to hydrogen bond.
The way that hydrogen bonding works is that, for example, it keeps water together. Water is here. And then what you wind up getting is that another water molecule gets close and between one of the O's and one of the H's you end up getting a loose – actually a pretty tight interaction. And that interaction is called hydrogen bonding.
Now this one's drawn a little bit weird, but you can see that pretty much every O can interact with an H and every H can interact with an O pretty much. They can all arrange themselves tightly together. That's what keeps molecules like water, like ethanol, that's what keeps them as liquids because they're able to get really close to each other and interact and attract each other.
Hydrogen bonding is going to be the most important intermolecular force. If any molecule has this force, we're going to say that that one is going to have the highest boiling point or the highest melting point or whatever. 

So now let's go to the second most important, the second strongest, and that's the thing we call the dipole-dipole force or what I like to call the net dipole force. Why? Because if you know how to draw dipoles and if you know how to find net dipoles, that's all you need for this.
For example, a molecule like acetone, which is one that I brought up earlier in our lessons, acetone looks like this. Does it have a dipole? Yes, it does. It actually has several dipole moments. It would have dipole moments pulling this way and this way because of the lone pairs. And it would have a major dipole moment pulling that way because of the double bond, but overall we would just say that it has basically a partial negative up here and a partial positive down here.
Does that make sense so far? Because basically, what's going to end up happening is that I'm just going to get a big dipole pulling electrons, oops. I'm going to get a big dipole pulling electrons towards the O. So I'm going to get a partial negative and a partial positive. Remember that we use the lower case delta to represent partial and it means that I don't know exactly what the number is, I just know that it's more than the other.
Well, check out what can happen. When I have a net dipole, another acetone molecule can arrange itself so that the partial negative from one of the molecules orients with the partial positive of the other and in that way they wind up sticking together. And you can imagine that I could have a bunch of acetone molecules all neatly arranged so that all the negatives link with all the positives. This is only possible though if you have a net dipole. If there's no net dipole, then you're not going to be able to form this force. Does that make sense?
So this is what we call the second strongest force. It's not quite as strong as hydrogen bonding, but it still is a pretty strong force. 

So then we move on to our last one which is our weakest force and it turns out that all molecules contain this force. Oops. All molecules possess Van der Waals. So I'm just going to put Van der Waals. But they don't all possess it to the same extent. What that means is that some molecules are going to have higher Van der Waals and some are going to have lower.
What makes Van der Waals' forces increase? The very first thing and most important thing is the size and that has to do with the molecular weight of the molecule. The higher the molecular weight, the stronger the Van der Waals. Pretty easy, right? Cool.
Then we've got the second most important thing or the second indicator is going to be the shape, the shape of the molecule. The shape has to do with how neatly they can be arranged and how neatly they can be ordered.
So what I'm going to do here is I'm going to draw three different molecules and you tell me which one would have the highest Van der Waals force. Here I've got a ring that's a six-membered ring. Here I've got a six-membered chain and then here I've got another six-membered chain. What I want to know is out of these three, which of them is going to have the highest Van der Waals.
First of all what we would look at is the size. Are all the sizes the same? Yes, they all have six carbons, so actually, in terms of size, they're fine.
Then I would look at the shape. Which one has the shape that can arrange the neatest and can stack the best? And the answer is the ring because check it out. The ring, since it's so symmetrical, it can have a bunch of rings stacked on top of each other. So if I wanted, I could keep drawing these rings stacked on top of each other. Do you know what that means? That there's intermolecular Van der Waals' forces all in between that keeping them stuck together. Does that make sense so far?
Then with the chains, the chains are still pretty good, but they're not quite as good as the ring. So here I'm drawing an example of Van der Waals' forces here. These are okay, but they're not quite as good as the rings. Does that make sense?
Then finally, you've got the branch, which I'll try to move out of the way because I know that it's right on top of me. But with the branch what we find is that there's not really a great way for them to stack together. They wind up kind of having a lot of space in between them. Does that make sense? So the Van der Waals' forces here are going to be very small compared to the Van der Waals' forces here for the rings. Does that make sense? Cool.
Basically, I hope that this clarifies intermolecular forces. The first thing to look for is the type of force. Hydrogen bonding is the best, Van der Waals is the worst, dipole is in the middle. And then we kind of break it down from there.
So now what I want to do is do a few practice problems so you guys get comfortable with this and I'm going to try and trick you a little bit, so be prepared. All right? Let's go.