- London dispersion forces
- Dipole–dipole forces
- Hydrogen bonding
- Ion–dipole forces
- Intermolecular forces and vapor pressure
- Solubility and intermolecular forces
- Surface tension
- Capillary action and why we see a meniscus
- Boiling points of organic compounds
- Boiling point comparison: AP Chemistry multiple choice
- Solubility of organic compounds
- 2015 AP Chemistry free response 2f
- Intermolecular forces
- Intermolecular forces and properties of liquids
Intermolecular forces and vapor pressure
A liquid’s vapor pressure is directly related to the intermolecular forces present between its molecules. The stronger these forces, the lower the rate of evaporation and the lower the vapor pressure. Created by Sal Khan.
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- do gases exert pressure equally in all directons?(4 votes)
- Yes they do, due to the random movements of the gas particles in all directions. It's important to remember, however, that not all of the particles have the same velocities. Some are faster than others, which is why pressure is the average of all of the forces exerted on the surfaces by the gas particles.(7 votes)
- At the beginning of the video, when Sal was figuring out the boiling points, he was looking at the O-H bonds. I thought these were intramolecular forces because they are within the molecule. Why exactly would it be intermolecular forces?(4 votes)
- So Sal is looking at the -OH groups on each of the molecules, but with the purpose of comparing their hydrogen bonding to other molecules which is of course an intermolecular force. When he's looking at the -OH bonds for each molecule, he's imagining other molecules of the same chemical in close proximity engaging in hydrogen bonding thus creating stronger attractions between the molecules which affects their boiling points. This is also why he investigates their London dispersion forces which is weaker intermolecular force.
Hope that helps.(6 votes)
- The types of intermolecular forces in a substance are identical whether it is a solid, a liquid, or a gas. Why then does a substance change phase from a gas to a liquid or to a solid?(2 votes)
- The physical states differ in how close the particles of matter are to each other essentially. With solid's particles being closely packed and having little motion relative to each other, and gas particles being greatly disperse from each other. These physical states also differ in the amount of kinetic energy the particles have, with gases having the most and solids having the least. Changing physical states requires the lose or addition of energy for matter then. The amount and strength of intermolecular forces tells essentially how much energy we need to change physical states.
If we consider water for example, changing solid water (ice) to liquid water is a physical change where we need to overcome the intermolecular forces of the water molecules and separate them. The amount of energy we need to add to make this physical change must overcome all the intermolecular forces water possesses. Any less than that and we simply don't have a physical change.
So physical changes still occur like we've learned in basic science classes, it's just that intermolecular forces provide barriers to those changes.
Hope that helps.(2 votes)
- is Methanol the same thing as Breaking bad? just kidding. Would like to know which video tells us about Dipole-Dipole forces, Thanks.(1 vote)
- Here's KA's video on dipole-dipole: https://www.khanacademy.org/science/ap-chemistry-beta/x2eef969c74e0d802:intermolecular-forces-and-properties/x2eef969c74e0d802:intermolecular-forces/v/dipole-dipole-forces
As for methanol and breaking bad, they did use it in the series. It shouldn't be confused with a number of other chemicals they used which also had the meth- prefix. In chemistry (particularly organic chemistry) chemicals with meth- in their names are common which stands for a methyl group (derived from the methane molecule) which is just a -CH3 group.
So methanol is a type of alcohol with a methyl group. It's commonly used in reactions as a solvent (something which other chemicals dissolve into it and hosts a reaction). In the show Walter uses methanol in the extraction of pseudoephedrine from cold medicine.
The show revolves around them synthesizing N-methylamphetamine, or commonly referred to as meth. Again, the meth just means there's a methyl group bonded to the rest of the molecule (the amphetamine).
Later in the series they also use methylamine as a different route to synthesize meth which doesn't require pseudoephedrine. Here we have a methyl group bonded to an amine (think of ammonia with a -CH3 group instead of one of the hydrogens).
Hope that helps.(4 votes)
- Does the term "hydrogen bond" refer to the bond between the hydrogen and an atom in its own molecule or the attraction between the hydrogen and another molecule?
For example, in the case of a mix of H2O and some negative dipole, is the hydrogen bond referring to the bond between H & O, or is it referring to the attraction between H and the negative dipole of another molecule?
sorry if that was a little bit wordy(1 vote)
- Hydrogen bonding is an intermolecular force, so it acts between molecules. As contrasted with an intramolecular force which acts within a molecule. Hydrogen bonding is just a stronger instance of dipole-dipole where the hydrogen of one molecule is attracted to an electronegative atom in another molecule. It's the same mechanism, it's just that hydrogen bonding only generally applies to molecules where hydrogen is directly bonded to fluorine, oxygen, or nitrogen.
Hope that helps.(3 votes)
- [Instructor] So we have four different molecules here. And what I want you to think about, if you had a pure sample of each, which of those pure samples would have the highest boiling point, second highest, third highest, and fourth highest? Pause this video, and try to figure that out. All right, now to figure that out, it really just boils down to which of these has the highest intermolecular forces when they're in a liquid state? Because if you have high intermolecular forces, it would take a lot of energy or a higher boiling point to really overcome those intermolecular forces and get to a gas state. So let's think about the intermolecular forces that we have studied. So I will start with hydrogen bonds, hydrogen bonds. 'Cause you could really view those, those are the strongest of the dipole-dipole interactions, and they're going to be stronger than your London dispersion forces. We can see that diethyl ether won't form hydrogen bonds. We don't see any bonds between hydrogen and an oxygen, a nitrogen, or a fluorine. Ethanol has one oxygen-hydrogen bond. Methanol also has one oxygen-hydrogen bond. Water has two oxygen-hydrogen bonds. So if I had to rank the hydrogen bond contribution to the intermolecular forces, I would put water as number one 'cause it can form the most hydrogen bonds. I would put methanol and ethanol as a tie for second. And then I would put diethyl ether last 'cause it can't form hydrogen bonds. So just looking at this, I know that water's going to have the highest boiling point. Diethyl ether is going to have the lowest boiling point. But what about the difference between methanol and ethanol? And we could think about other types of dipole forces, but not a lot that you could intuit just by eyeballing them. They might actually have similar dipole moments on a molecular basis. But we can think about London dispersion forces. I'll do this in a different color. So London dispersion forces. And if we're just trying to, actually I'll rank all of them. So London dispersion forces are proportional to how polarizable a molecule is, which is proportional to how large its electron cloud is, which is proportional to its molar mass. And it's clear that diethyl ether has the highest molar mass, followed by ethanol, followed by methanol, followed by water. How did I know that? Well, you literally can take atoms away from the diethyl ether to get to an ethanol. And you can literally take atoms away from that to get to a methanol. And you can literally take atoms away from that to get to a water. So we know that this is the order of molar mass. And so London dispersion forces, I wouldn't make that change the ranking between water or diethyl ether because these are going to be a lot weaker than those hydrogen bonds. But they can be useful for the tiebreaker between ethanol and methanol. And so my overall ranking on boiling points, the highest boiling point I would put would be water, followed by, since ethanol won the tiebreaker, followed by ethanol, followed by methanol, and then the lowest boiling point would be diethyl ether. And if we look at the actual data, it's consistent with what we just talked about. We can see very clearly that water has the highest boiling point, ethanol is second, methanol is third, and diethyl ether was fourth, completely consistent with our intuition. Now, what's also interesting here, you might have noticed, is this thing called vapor pressure. And you might have also noticed that vapor pressure seems to trend the opposite way as boiling point. The things that have the high boiling point have the low vapor pressure, and the things that have the low boiling point have a high vapor pressure. So what are we talking about, why, about vapor pressure, and why do we see this relationship? And I'm not going to go deep into vapor pressure. There'll be other videos on that on Khan Academy. But just to get you a sense, imagine a closed container here. And I put one of these, a sample of one of these molecules in a liquid state, and I'm gonna just draw the molecules, clearly not drawn to scale, as these little circles. And the temperature matters, so let's say that this is at 20 degrees Celsius. Now, you might notice, at 20 degrees Celsius, it's lower than the boiling point of all of these characters. So for the most part, they're going to be in a liquid state, but we know that not every one of these molecules is moving with the exact same kinetic energy. The temperature, you could view as a measure of the average kinetic energy of the molecules, but they're all bumping around into each other, in different positions, with different amounts of velocities and therefore different kinetic energies. And so every now and then, you're going to have a molecule that has the right position and the right kinetic energy to escape and get into the vapor state, into a gaseous state. And so that's going to keep happening. But then the things that are in the gaseous state, every now and then they're bumping into each other, and they're bumping into the sides of the container. And every now and then, they might approach the surface with the right kinetic energy, with the right position, so that they get recaptured by the intermolecular forces and enter a liquid state. And so you can imagine, this will keep happening where things go from liquid, and then they go to vapor. But then when that vapor gets high enough or when you could say the vapor pressure gets high enough, remember, that pressure's just from the vapor molecules bouncing around, then you will get to some form of an equilibrium. And you could imagine, the things that have a lower boiling point, that means they have lower intermolecular forces, more of the vapor is going to form, and so you're going to have a higher vapor pressure before you get to equilibrium. On the other hand, things with high intermolecular forces, fewer of those molecules are going to break away, and so you're going to have a lower vapor pressure when you get to that equilibrium. And you can see that very clearly here. So I will leave you there. We got a little bit of practice, seeing everything we've seen so far, and we learned a little bit about vapor pressure and how that relates to intermolecular forces and boiling point.