more the intermolecular forces, lesser is the vapour pressure this is how it goes Vapor pressure is In the end, we can make the overall relationship that: . As the intermolecular attraction increases,. • The vapor pressure (the pressure of the vapor that is in equilibrium with its liquid) decreases. • The boiling point. Answer to What is the relationship between intermolecular forces in a liquid and the liquid's boiling point and critical.
Which implies, and it's true, that all of the molecules do not have the same kinetic energy. Let's say even they did. Then these guys would bump into this guy, and you could think of them as billiard balls, and they transfer all of the momentum to this guy.
Now this guy has a ton of kinetic energy. These guys have a lot less. This guy has a ton. There's a huge distribution of kinetic energy. If you look at the surface atoms or the surface molecules, and I care about the surface molecules because those are the first ones to vaporize or-- I shouldn't jump the gun. They're the ones capable of leaving if they had enough kinetic energy. If I were to draw a distribution of the surface molecules-- let me draw a little graph here.
So in this dimension, I have kinetic energy, and on this dimension, this is just a relative concentration. And this is just my best estimate, but it should give you the idea.
So there's some average kinetic energy at some temperature, right? This is the average kinetic energy. And then the kinetic energy of all the parts, it's going to be a distribution around that, so maybe it looks something like this: You could watch the statistics videos to learn more about the normal distribution, but I think the normal distribution-- this is supposed to be a normal, so it just gets smaller and smaller as you go there.
And so at any given time, although the average is here, there's some molecules that have a very low kinetic energy. They're moving slowly or maybe they have-- well, let's just say they're moving slowly. And at any given time, you have some molecules that have a very high kinetic energy, maybe just because of the random bumps that it gets from other molecules.
It's accrued a lot of velocity or at least a lot of momentum. So the question arises, are any of these molecules fast enough? Do they have enough kinetic energy to escape? And so there is some kinetic energy. I'll draw some threshold here, where if you have more than that amount of kinetic energy, you actually have enough to escape if you are surface atom. Now, there could be a dude down here who has a ton of kinetic energy. But in order for him to escape, he'd have to bump through all these other liquid molecules on the way out, so it's a very-- in fact, he probably won't escape.
It's the surface atoms that we care about because those are the ones that are interfacing directly with the pressure outside. So let's say this is the gas outside. It's going to be much less dense. It doesn't have to be, but let's assume it is.
These are the guys that kind of can escape into the air above it, if we assume that there's some air above it. So at any given time, there's some fraction of the particles or the molecules that can escape.
So you're next question is, hey, well, doesn't that mean that they will be vaporized or they will turn into gas? And yes, it does. So at any given time, you have some molecules that are escaping.
Those molecules-- what it's called is evaporation. This isn't a foreign concept to you. If you leave water outside, it will evaporate, even though outside, hopefully, in your place, is below the boiling temperature, or the normal boiling temperature of water. The normal boiling point is just the boiling point at atmospheric pressure.
If you just leave water out, over time, it will evaporate.
What happens is some of these molecules that have unusually high kinetic energy do escape. They do escape, and if you have your pot or pan outside or, even better, outside of your house, what happens is they escape, and then the wind blows. The wind will blow and then blow these guys away. And then a few more will escape, the wind blows and blows them all away.
And a few more escape, and the wind blows and blows them all the way. So over time, you'll end up with an empty pan that once held water. Now, the question is what happens if you have a closed system? Well, we've all done that experiment, either on purpose or inadvertently, leaving something outside and seeing that the water will evaporate. What happens in a closed system where there isn't wind to blow away? So let me just draw-- there you go. Let's say a closed system, and I have-- it doesn't have to be water, but I have some liquid down here.
And there's some pressure from the air above it. Let's just say it was at atmospheric pressure.
Vapor Pressure ( Read ) | Chemistry | CK Foundation
It doesn't have to be. So there's some air and the air has some kinetic energy over here. So, of course, do the water molecules. And some of them start to evaporate.
So some of the water molecules that are up here in the distribution, they have enough energy to escape, so they start hanging out with the air molecules, right? Now something interesting happens. This is the distribution of the molecules in the liquid state. Well, there's also a distribution of the kinetic energies of the molecules in the gaseous state. Just like different things are bumping into each other and gaining and losing kinetic energy down here, the same thing is happening up here.
So maybe this guy has a lot of kinetic energy, but he bumps into stuff and he loses it. And then he'll come back down. So there's some set of molecules.
I'll do it in another set of blue. These are still the water-- or whatever the fluid we're talking about-- that come back from the vapor state back into the liquid state. And so what happens is, there's always a bit of evaporation and there's always a bit of condensation because you always have this distribution of kinetic energies.
At any given moment in time, out of the vapor above the liquid, some of the vapor loses its kinetic energy and then it goes back into the liquid state. Some of the surface liquid gains kinetic energy by random bumps and whatever else and goes into the vapor state. And the vapor state will continue to happen until you get to some type of equilibrium. And when you get that equilibrium, we're at some pressure up here. So let me see, some pressure. And the pressure is caused by these vapor particles over here, and that pressure is called the vapor pressure.
I want to make sure you understand this. So the vapor pressure is the pressure created, and this is at a given temperature for a given molecule, right? Every molecule or every type of substance will have a different vapor pressure at different temperatures, and obviously every different type of substance will also have different vapor pressures.
For a given temperature and a given molecule, it's the pressure at which you have a pressure created by the vapor molecules where you have an equilibrium. Where you have just as many things vaporizing as things going back into the liquid state. And we learned before that the more pressure you have, the harder it is to vaporize even more, right?
We learned in the phase state things that if you are at degrees at ultra-high pressure, and you were dealing with water, you would still be in the liquid state. So the vapor creates some pressure and it'll keep happening, depending on how badly this liquid wants to evaporate. But it keeps vaporizing until the point that you have just as much-- I guess you could kind of view it as density up here, but I don't want to think-- you have just as many molecules here converting into this state as molecules here converting into this state.
So just to get an intuition of what vapor pressure is or how it goes with different molecules, molecules that really want to evaporate-- and so why would a molecule want to evaporate? It could have high kinetic energy, so this would be at a high temperature. It could have low intermolecular forces, right? It could be molecular. Obviously, the noble gases have very low molecular forces, but in general, most hydrocarbons or gasoline or methane or all of these things, they really want to evaporate because they have much lower intermolecular forces than, say, water.
Or they could just be light molecules. You could look at the physics lectures, but kinetic energy it's a function of mass and velocity. So you could have a pretty respectable kinetic energy because you have a high mass and a low velocity. This would result in a slight attraction of the two molecules until the charges moved around again but is responsible for the attractive London dispersion forces all molecules have.
However, these London dispersion forces are weak, the weakest of all the intermolecular forces.
Their strength increases with increasing total electrons. Dipole-dipole attractions What would happen if we had a beaker of polar molecules, like formaldehyde, In addition to the attractive London dispersion forces, we now have a situation where the molecule is polar.
We say that the molecule has a permanent dipole. Now, the molecules line up. The positive ends end up near to another molecule's negative end: Since this dipole is permanent, the attraction is stronger.
However, we only see this sort of attraction between molecules that are polar. It is usually referred to as dipole - dipole interaction.
The strength of this attraction increases with increasing total number of electrons. Hydrogen bond Hydrogen is a special element. Because it is really just a proton, it turns out that it can form a special type intermolecular interaction called the hydrogen bond. If the hydrogen in a moleucle is bonded to a highly electronegative atom in the second row only N, O, or Fa hydrogen bond will be formed. In essence the three elements listed above will grab the electrons for itself, and leave the hydrogen atom with virtually no electron density since it had only the one.Which molecules have higher (or lower) vapor pressure
Now, if another molecule comes along with a lone pair, the hydrogen will try to position itself near that lone pair in order to get some electron density back. This ends up forming a partial bond, which we describe as the hydrogen bond. The strength of this interaction, while not quite as strong as a covalent bond, is the strongest of all the intermolecular forces except for the ionic bond.
A diagram of the hydrogen bond is here: Could the CH2O molecule exhibit hydrogen bonding? The answer is no, since the hydrogen must be bound to either N, O, or F. Just having one of those species in the molecule is not enough. Trends in the forces While the intramolecular forces keep the atoms in a moleucle together and are the basis for the chemical properties, the intermolecular forces are those that keep the molecules themselves together and are virtually responsible for all the physical properties of a material.
The intermolecular forces increase in strength according to the following: Therefore, one would expect the melting and boiling points to be higher for those substances which have strong intermolecular forces. We know that it takes energy to go from a solid to a liquid to a gas. This energy is directly related to the strength of attraction between molecules in the condensed phases. Since energy is directly proportional to the temperature, the above trends ought to hold true.
In addition, there are energies associated with making these phase transitions: Each of these processes are endothermic, and scale with the magnitude of the intermolecular forces.