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Current time:0:00Total duration:18:04

Video transcript

we know that when we have some substance in a liquid state it has enough kinetic energy for the molecules to move past each other but still not enough energy for the molecules to completely move away from each other so for example this is the liquid maybe that they're moving in that direction these guys are moving a little bit slower in that direction so there's a bit of this flow going on but still their bonds between them they kind of switch between different molecules but they want to stay close to each other there are there are these little bonds between them and they want to stay close and then if you if you increase the coverage kinetic energy enough or essentially increase the temperature enough and then overcome the heat of fusion we know that all of a sudden even these bonds aren't strong enough to even keep them close and the molecules separate and begin to into a gaseous phase and there they have a lot of kinetic energy and they're bouncing around and they take the shape of their container but there's an interesting thing to think about temperature is average kinetic energy temperature is average kinetic energy which implies and it's true that all of the molecules do not have the same kinetic energy that you know maybe let's say even they did then these guys would bump into this guy and you know the Coulomb is billiard balls and they transfer all of the momentum to this guy now this guy has a ton of kinetic energy this guy these guys have a lot less this guy has a ton these guys have a lot less there's a huge distribution of kinetic energy if you look at the surface atoms and or the surface molecules and I care about the surface molecules because those are the first ones to vaporize or or I shouldn't I shouldn't jump the gun they're the they're that 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 on this dimension I have kinetic energy kinetic energy and on this dimension this is just the relative concentration and this just you know my best estimate but it should get you the give you the idea so there's some average kinetic energy at some temperature all right this is the average kinetic energy and then the kinetic energy of all the but it's going to be a distribution around that so maybe it looks something like this a bell curve and 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's just it's smaller and smaller as you go there and so at any given time although the averages here there are some molecules that have a very low kinetic energy they're moving slowly or maybe they have well let's say say they're moving slowly and then at any given time you have some kit some modules that have a very high kinetic energy maybe just because of the random bumps and and and it gets from other molecules it's a it's a crude a lot of velocity or at least a lot of momentum so the question arises are any of these molecules fast enough are they do they have enough kinetic energy to escape and so there is some kinetic energy I'll draw some threshold here or if you have a more than that amount of kinetic energy you actually have enough to escape if you are a surface atom now there could be there could be a dude down here who has a ton of kinetic energy but he's going to in order for him to escape you'd have to bump through all of 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 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 alright so 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 of the molecules that can escape so your next question is hey well doesn't that mean that they will vaporize or they will turn into gas and yes it does so at any given time you have some molecules that are escaping with those molecules what it's called is evaporation and 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 so some of these molecules what happens is some of these molecules that have unusually high kinetic energy do escape they do escape and if you have your your pot or pan outside or even better outside of your house then what happens 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 blow them all away and a few more escape and the wind blows and blows them all away so over time you'll end up with a with an empty pan that once held water now the question is what happens if you have a closed system well we know I mean we've all done that experiment you either on purpose or inadvertently leaving something outside and seeing that the water will will evaporate what happens in a closed system where there isn't wind to blow away so let's let me just draw nope - there you go say closed system and I have it doesn't have to be water but I have some liquid down here some liquid down here and there are some pressure from the air above it let's just say it was that atmospheric pressure doesn't have to be so there's some air and the air has some kinetic energy over here and so of course do the water molecules and some of them start to evaporate so some of the water molecules that have you know 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 of the kinetic energies of the molecule and the gaseous state just like you know 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 you know there's some set of molecules I do 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 this this evaporation well there's always a bit of evaporation and there's always a bit of condensation because you always have this distribution of kinetic energies there's some 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 in the liquid state some of the surface liquid gains kinetic energy by random bumps and whatever else and goes into the vapor State and and that'll the vapor State will continue to happen until you get to some type of equilibrium and when you get that equilibrium at some pressure up here so let me see some pressure and the pressure is caused by these vapor particles over here these vapor particles right here and that pressure is called the vapor pressure vapor pressure I want to make sure you understand this so the vapor pressure is the pressure created at and this is at a given temperature at a given temperature for a given molecule given temperature for a given molecule right every every molecule or every type of substance will have a different vapor pressure different temperatures and obviously every different type of substance will also have different vapor pressures before a given temperature in a given molecules the pressure at which you have an equal a pressure created by the vapor molecules where you have an equilibrium we 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 put if you if you are a 100 degrees at ultra-high pressure you would still and you were dealing with water you would still be in the liquid state so this 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 that 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 in 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 so want to evaporate want to evaporate and so why would a molecule want to evaporate it could have high high kinetic energy so this would be at a high temperature it could have low intermolecular forces all right it could be molecular obviously like the noble gases have very low molecular forces but in general most most hydrocarbons or you know gasoline or methane or all of these things they really want to evaporate because they have much lower intermolecular forces and say water or they could just be light molecules light molecules you could 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 but if you because you have a high mass and a low velocity so if you have a if you have a light mass and the same kinetic energy or more likely to have a higher velocity you could watch the kinetic energy videos for that but something that wants to evaporate it's a lot of it's a lot of its molecules let me do it in a different color something that wants to evaporate really bad a lot more of its molecules will have to enter into this vapor State in order for the equilibrium to be reached let me do it all in the same color so the pressure created by its vaporize its evaporated molecules is going to be higher for to get to that equilibrium state so it has high vapor pressure so high vapor pressure and on the other side if you're at a low temperature or you have strong intermolecular forces or you have a heavy molecule or I then you're going to have a low vapor pressure for example like you know I mean iron has a very low vapor pressure because it's not vaporizing while I'll let me think of something you know carbon dioxide has a relatively much higher vapor pressure much more of carbon dioxide is going to evaporate when you have it well I really shouldn't use that because you're going straight from the liquid to the solid state but I think you get the idea and something that has a high vapor pressure that wants to evaporate really bad we call it we say it has a high volatility you've probably heard that word before high volatility high volatility so for example gasoline has a higher it's more volatile than water and that's why it evaporates and it also has a higher vapor pressure because if you were to put it in a closed container more gasoline at the same temperature will enter and the same atmospheric pressure will enter into the vapor state and so that vapor state there's more the vapor state will generate more pressure to offset what we'll have to generate more pressure to offset the natural inclination of the gasoline to want to escape than in the case with water now an interesting thing happens when this when this vapor pressure is equal to the atmospheric pressure so right now this was our closed container and we had some you know you have the atmosphere here at a certain pressure and let's say until now we've assumed that the atmosphere the atmosphere was at a higher pressure so it was for the most part keeping these molecules contained maybe some atmosphere molecules are coming in here and maybe some of the vapor molecules or escaping a bit but it's keeping it contained because this is at a higher pressure out here than this vapor pressure and of course the pressure right right here at the surface of the molecule it's going to be the combination of the partial pressure due to the to the few atmospheric molecules that come in plus the vapor pressure but once that vapor pressure becomes equal to that atmospheric pressure so it starts to it can press out with the same amount of force you can kind of view it or the force per area so then the molecules can start to escape I can push the atmosphere back and so you start having a gap here you start having you start having a vacuum I would you know I don't to use exactly a vacuum but it's in some oxygen escape you more and more of these molecules can start going out and at that point you've reached the boiling point of the substance when the vapor pressure is equal to the atmospheric pressure now just to get a sense of what all some of this means let's look at the vapor pressure for water this is water right here h2o I should do that in black h2o and so you see at at 760 so atmospheric pressure we're in tours now but that's just a different one 760 tours is equal to one atmosphere so that's about right number one that's about right there so it's one atmosphere so at atmospheric pressure the vapor pressure at a hundred degrees Celsius for water the vapor is is at 100 degree Celsius for water or I guess another way to put at 100 degrees Celsius you have 760 Torr a vapor pressure which is exactly the the atmospheric pressure at a one atmosphere at sea level so at 100 degrees vapor pressure is equal to atmospheric or sea level atmospheric sea level and so you're going to boil which we all know is true and then if you have at lower temperatures at lower temperatures your vapor pressure your vapor pressure is going to be lower than the atmospheric pressure right let's see it look here it looks like 300 something but then what happens if you lowered the atmospheric pressure enough if you were to pump air out of the the container or whatever low enough and then the so if you brought the atmospheric pressure down to this vapor pressure then again you will have boiling and we saw that in the phase change diagrams that you can boil something at a lower temperature if you lower the atmospheric pressure and that's because you're lowering the atmospheric pressure to the vapor pressure of the substance and here's a comparative chart of a and this is interesting you see this is kind of a exponential increase with temperature of vapor pressure and that's because if you think about that distribution we did before this is at one kinetic energy if you increase the amount of kinetic energy then your distribution will look like this if the temperature has gone up and now you have a lot lot more it's not just linear you have a lot more particles that can now escape and have the kinetic energy to evaporate and you can see it's this exponential increase as you increase the temperature now here is another chart you say where is that exponential increase go and that's because this is a logarithmic you could see the scale it goes on it increases with with exponentially as opposed to linearly so it goes from point 1 to 10 so equal distances are actually up by a factor of 10 so that's why you don't see that logarithmic move but these are just four different substances propane you see at any given so let's go for let's go like a decent temperature let's go 20 degrees Celsius 20 degrees Celsius propane has the highest vapor pressure so this is one atmosphere so propane will actually will actually evaporate will actually boil at 20 degrees Celsius it'll actually completely boil and go into the gaseous state because it's vapor pressure is so much higher than atmospheric pressure if we're assuming it's we're at sea level and then you could do that for different different molecules methyl chloride is the next one slightly lower vapor pressure but still very volatile it would still definitely boil and turn into the gaseous state at twenty degrees Celsius if we're at sea level because sea level is right there and let's see it at sea level if you wanted to keep something so sea level is this pressure if you wanted to keep if you wanted to keep let's say methyl chloride if you wanted to keep methyl chloride in the liquid state or in equilibrium with the liquid state instead of boiling you would have to be at least at around what is this - 25 degrees Celsius in order for that even a propane even at minus 25 degrees it's still in the gaseous state because it's vapor vapor pressure is still higher and then of course that if you go onto the you know if you butane for example butane I think is what they put actually butane is what they put in I think they put it in lighters but butane will be in the liquid state as long as you're at around roughly 0 degrees and in a lighter you might tell it's in a liquid state they probably increase the pressure so the pressure and the lighter is probably something higher maybe it's maybe it's at 2 atmospheres or something so that the butane is at room temperature will stay in the liquid state who knows I don't what the pressure is in there there's just an interesting chart to look at that there's actually a bunch of different vapor pressures you can see at atmospheric pressure what's likely to be a gas or liquid in temperatures and then you could see at different temperatures which are the things that are most volatile and how much do you have to increase or decrease the pressure to to evaporate something or to boil it anyway hopefully you found that useful vapor pressure it's something that we encounter every day and I'll see you in the next video
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