in this lesson on liquids and solids we're gonna have a pretty comprehensive discussion about intermolecular forces including hydrogen bonding dipole-dipole forces London dispersion forces ion dipole forces well you see how those relate to the bulk properties of like boiling point melting point surface tension viscosity vapor pressure and then we'll have a discussion about phase diagrams a plot of pressure versus temperature where solids liquid gases are we'll talk about the lines of equilibrium the critical point the triple point things of this sort and if you learn anything from this lesson liked it share it and subscribe if you'd like to be notified about future lessons I release and if you're looking for good study guides and practice problems check out Chad's prep comm my ultimate general chemistry prep course has currently over 1200 questions and Counting all right so in a discussion of intermolecular forces here grana first start off by limiting ourselves to talking about these three intermolecular forces hydrogen bonding dipole-dipole forces and leonard as written forces and the reason we'll start with just these three is that these three are the only options for a pure liquid in a pure liquid this is all that's available for the interaction between individual molecules now when we get to mixtures we'll add on ion dipole forces onto this now these three can also all be present mixtures but I indict phille forces imply that you have to have a mixture which is why we'll leave them out of the discussion for now let's talk about first what an intermolecular force is once then we take a look at a couple of HCl molecules here so and my first question for you is in a lewis structure here what is that horizontal line depict right there and hopefully you said covalent bond because that would be correct and a covalent bond is not at all we're talking about when we talk about intermolecular forces inter molecular implies between separate molecules so it turns out the covalent bond actually holds a single molecule HCL together and in that case we might call it an intra molecular force so things of a sort but the force we're talking about in this lesson is the force between separate molecules and it turns out there's a rather weak attractive force much weaker than an actual covalent bond so that's actually attracting this too and it's just due to plus and minus in every one of these cases so if we look a CH CL is a polar molecule so and chlorine being more electronegative is partially negative and hydrogen being less electronegative is partially positive and on the adjacent molecule it is exactly the same situation and so the traction here is due to the fact that the chlorine on one molecule is partially negative and is attracted to the partially positive hydrogen on the next molecule over now the nature of this force is due to the fact that this molecule on the left is polar and this molecule on the right is polar and obviously they're identical molecules in a sample of HCl you're gonna have zillions of molecules I'm just examining two so I'm showing this in a pure liquid let's say but this would also happen in a mixture with different components as long as they were both polar so because they're both polar they both have a dipole moment and so this interaction between two polar molecules is called a dipole-dipole force so I'm just I'm not starting at the top I'm kind of starting in the middle here and this is of intermediate strength we're gonna find they get stronger as we go up with hydrogen bonding being the strongest of these three cools this is a dipole-dipole force and you just have to have polar molecules in this case to happen it's just partial positive attraction to partial negative attraction on another molecule now hydrogen bonding I like to think of as a super duper strong dipole-dipole force now I don't particularly care for the Dame one we use the word bonding in the name and it makes students they go it's kind of as strong as a covalent bond right no this is just an intermolecular force and again it's significantly significantly weaker than an actual covalent bond so let's keep that straight also it's called hydrogen bonding and it makes students think that all you got to have is hydrogen and you can do hydrogen bonding well that's false - there's only three types of hydrogen atoms that can be involved in hydrogen bonding and that is a hydrogen bonded to a fluorine a hydrogen bonded to an oxygen or a hydrogen bonded to a nitrogen so I like to call these the phone elements F Oh n-no sir the only hydrogens that can be involved in the idea is that fluorine is the most electronegative elements oxygen is second and nitrogens in the running for third or fourth depending on who you talk to so but with these three very polar bonds you're gonna get a hydrogen that has a very significant amount of partial positive charge and it's only those hydrogens that can actually be involved in hydrogen bonding so like the hydrogen bonded to a chlorine it's polar it's partially positive but it's not as partially positive as these guys and it's not gonna be involved in hydrogen bonding so with one of these three very very polar bonds to hydrogen you can have hydrogen bonding and water is very well known for having a significant degree of hydrogen bonding and so if we take a look at a couple of different water molecules here so the Hydra money we'll look at is the interaction between a very partially positive hydrogen and one molecule and the partially negative oxygen of the adjacent molecule and technically it's actually mediated through the lone pair of electrons from the oxygen and so in this case with hydrogen bonding would call the molecule that has the hydrogen involved the hydrogen bond donor we call the molecule that has the fluorine oxygen or negative sorry fluorine oxygen or nitrogen with a lone pair involved the hydrogen bond acceptor and so in this case each water molecule having two H's bonded to o can act as a hydrogen bond or up to twice so and having two lone pairs on the oxygen can act as a hydrogen bond acceptor up to twice and so each water molecule can actually be involved in hydrogen bonding with four other water molecules now in the liquid phase you know usually each water molecule you know they're moving around and stuff is on average hydrogen bonding with somewhere between two and three other molecules but when you freeze water so it expands into ice and it turns out it expands so that every single water molecule fits into a crystal structure where each water molecule is interacting with four other water molecules through hydrogen bonding and so that's what causes water to expand is this hydrogen bonding giving it a very favorable structure with this very strong intermolecular force involved all right so hydron bonding again like a super duper duper strong dipole-dipole force and generally considerably stronger than all the other regular dipole-dipole forces which is why they just decided to give it its own special name alright so we'll talk about London dispersion forces so in London dispersion forces it turns out that all molecules have them water has them HCl has them so however most notably we'll talk about Lemon dispersion forces for non-polar molecules because if you're nonpolar you don't have hydrogen bonding and you don't have dipole-dipole which means the only intermolecular force you'll have is one industry reinforces but keep in mind again all molecules have these with water we just often don't talk about it because if you got you know super glue why would you talk about the scotch tape so that's kind of the idea but keep in mind that all molecules have these now if you're not in polar this is all you got so and if we kind of take a look at example let's say I'm a nonpolar molecule and you're a nonpolar molecule and there should be there for no reason we should be attracted to each other except for the fact that our electrons are in motion let's say this marker here represents my electron cloud and it's rotating around me and freeze and for this exact instant in time my electrons are facing you and so facing you I appear to be a little bit negative and so what you do is you take your electrons and spin them around behind you that way facing me you're a little bit positive and I look at you and you're positive and you look at me and I'm negative we say hey we should hang out sometime and then our both of our electrons move like oh never mind so it is a weak temporary or transient dipole it's not a permanent dipole if you have a permanent dipole you got dipole type forces but if it's just a weak temporary or transient dipole so just due to the motion of electrons then this is what it amounts to and it's the weakest of the intermolecular forces now the truth is though it totally depends on the size of the molecule as well as the surface area so it turns out bigger molecules are going to have more electrons in therefore have a bigger London dispersion forces and it turns out even bigger atoms as you go down to a group not only do they have more electrons but that we say they're more polarizable they have squishy or electron clouds which can cause them to have greater temporary dipoles if you will and so size plays a role but also surface area so if I told you that I was gonna take care this old review sheet here and I was gonna coat it in say Elmer's glue on one side and then post it up on the board here and the question is would it stick well it probably would probably would now if I took that same piece of paper still coated on one side in Elmer's glue and punched up into a ball and stuck it against the board here so it might stick it might not but we can say this it is way less likely to stick and so it turns out that's the surface area part so molecules with a smaller surface area therefore have a smaller surface area to interact with other molecules and are gonna have smaller London dispersion forces overall so again London dispersion forces depend on size and surface area and the truth is for really big molecules these can actually get pretty large so we said they are the weakest of the intermolecular forces but they're additive depending on size and really big molecules can actually overcome you know a lack of polarity or even some cases extreme cases a lack of hydrogen bonding and we'll talk about how you kind of distinguish those differences okay so the reason identifying these intermolecular forces important is it relates to some of the bulk properties of a liquid so like the melting point or the boiling points what we call the surface tension or the viscosity as well as the vapor pressure and before we get into those five things though we just want to take a look at a comparison here so let's say we get three molecules here and I'm gonna pick methanol here and we'll compare him to methane and hydrogen and we're gonna compare these three molecules and my question for you is which of these three has the greatest intermolecular forces and cluing you in real quick you should focus on that OAH bond right there from the get-go and if you have an OHA on then you're gonna be capable of hydrogen bonding this is going to interact with adjacent molecules with the strongest of these three intermolecular forces cool so if I just said which of these has the highest intermolecular force you'd say ch3oh now if i said which of these has the weakest intermolecular forces so you'd go with hydrogen well i already know he's the strongest or he's out but ch4 and h2 both are nonpolar and the only animal air force they have is a lot of dispersion forces but it's totally size dependent ch4 is molecular weight is 16 h2s is 2 and ch4 has a much bigger surface area as well and so due to its smaller size and surface area h2 has a significantly lower intermolecular forces alone dispersion forces then methane here in this case and so he'd have the weakest overall intermolecular forces now the reason that's important is again these intermolecular forces rate to the bulk properties of a liquid so if you look like boiling a liquid so in the liquid phase the molecules are touching but when you boil it they get separated out by huge amounts of empty space and so if they're touching in the liquid and then they're separated by empty space in the gas you have to break apart any stickiness between the molecules you have to overcome any intermolecular forces they have and so the stronger those intermolecular forces the more energy it's going to take and the higher the boiling temperature and it works similarly for the melting point as well so when you go from solid to liquid you have to break some of those intermolecular forces and again they're stronger they are the more energy that's going to take as well and so higher intermolecular forces leads to a higher boiling point and a higher melting point and so if I said which of these three compounds has the highest boiling point or the highest melting point you'd say the one that had the strongest intermolecular forces ch3oh again being capable of hydrogen bonding now the other bulk properties relate this to our surface tension and then viscosity so surface tension if you kind of think of what if I told you I have a friend and my friend when he jumps off the diving board into the swimming pool he actually walks on water just walks on the water so one you'd call me crazy so but then I might tell you that my friend is a water bug so one you might be like well Chad you just don't parently have enough friends or you're not getting out much I will tell you I'm not getting out much we're all not getting out much right now we're all in isolation right so but a water bug is my friend in this example and that water bug is one light enough and two it spreads out its weight over a big broad surface area that it doesn't actually break through the webbing of hydrogen bonding going on on the surface of the water so that webbing the water bugs light enough not to push through now you were i jumping off a diving board well we're just fat and we're gonna plow right through so that hydrant bonding Network so but that water bug light enough spreading its weight out of amongst a broad enough surface area that it doesn't and that's a measure of surface tension that webbing at the surface of water another place you'll see this is if you like to fill up a glass and if you fill it up slowly with a small trickle of water you can actually fill that glass to where it bubbles up at the top and you fill it a little fuller than the actual glass and the reason you can fill it up fuller is that the bubble on top you can you have this network of hydrogen bonding going on that holds that bubble in place so to speak so that's surface tension and the greater the intermolecular forces the greater the surface tension again and so if I ask you which of these three has the highest surface tension you'd say ch3oh yet again having the highest intermolecular forces now the next is viscosity and you know kind of layman's term we think about viscosity just think well how thickly the liquid pours or something like this you know and we think of things like honey and molasses and motor oil and things of this sort and so the technical definition of viscosity of those deals with fluid flowing in straight lines and so straight layers if you will and if you have two layers of fluid viscosity is the friction between those layers and the reason for that friction is because of attractive forces between the molecules in those layers and if there's a lot of attractive forces there'll be a lot of friction and they'll want to flow together which is why viscous things tend to flow thickly if you will if that's even a word that's kind of the idea and so the greater the intermolecular forces yet again then the greater the attraction between the layers and the more viscous that lay liquids gonna be and so I said which of these three would be the most viscous were have the highest viscosity it would be ch3oh yet again now the last one we got to talk about is one's gonna be kind of counter on the trends and it's gonna be called a vapor pressure and let's get a good picture going here so we're done with these for now so let's say we've got a container so and we're gonna have a liquid in this container so and there's gonna be some rather large molecules here for the purpose of this illustration in that container so in those molecules are all having some intermolecular forces that are attracting them to each other and stuff like this and they're moving around and some are faster and some are slower there's this distribution of kinetic energies and stuff so but every once while one of them makes a beeline for the surface so and if it's moving fast enough it will actually have enough kinetic energy to jump out into the gas phase so however some molecules won't have enough kinetic energy and instead of having enough to jump out to the gas phase it's still gonna be attracted all these liquid molecules and the attraction big later would be greater than its kinetic energy so only the fastest molecules have enough to jump out into the gas phase up here and form a vapor and the measure of how many of the molecules made it up here is called the vapor pressure the more molecules you have up here's vapor the greater the vapor pressure the fewer the molecules you have upper is vapor the lower the vapor pressure and the idea is that if you're liquid molecules have greater intermolecular forces well then you're gonna have to have molecules moving even faster to jump up into the vapor phase so and at any given temperature you only have a certain fraction of molecules that are going to have that much kinetic energy and the stickier these get the smaller that fraction becomes and the fewer the molecules that actually make it up here and so this is the one thing that's gonna go down the highest intermolecular forces doesn't have the highest vapor pressure it now actually has the lowest vapor pressure instead and so again I could ask you which of these compounds has the highest melting point the highest boiling point the highest surface tension the highest viscosity at ch3oh every time but if I say which one's got the highest vapor pressure while it definitely in him he's got the lowest vapor pressure the highest vapor pressure would be associated with the molecule with the weakest intermolecular forces we said was gonna be age two out of these three oh so those are the things you kind of compare based on intermolecular forces and we would only make a comparison of those five bulk properties for pure liquids we're not gonna you know say well if you've got methanol mixed with water and you've got you know this over here we're sure you know we're not going to make those comparisons on boiling points and stuff like that now we might say what kind of intermolecular forces you have in a mixture and which ones are stronger and stuff like that but we're not going to relate those to bulk properties only for pure liquids and that's why we restricted ourselves to talking only about these three intermolecular forces because they're the ones that can be involved in pure liquids now again they also can be in mixtures but they're the only ones in Mick involved in pure liquids so the one other intermolecular force we want to talk about ion dipole forces and even the name here implies that we have a mixture we've got something ionic mixed with something with a dipole something polar so a most common example of this is when you dissolve like salt in water you've gotta salts an ionic compound dissolved in water a polar solvent in this case and so say dissolve that salt and you're gonna get a sodium ion and a chloride ion and they're both gonna be interacting with some water molecules now for the sodium ions they're gonna be interacting with the oxygen of the water molecules because the sodium is partially positive sodium has got a full positive charge but the oxygens are partially negative and so this interaction right here is what we call the ion dipole force now it turns out it depends on you know the charge of your ion whether it's plus 1 plus 2 plus 3 minus 1 minus 2 minus 3 that that's gonna affect the strength and also the size of the ion turns out smaller ions will form stronger interactions and larger ions of form weaker ones and stuff like that but in general you should know that these ion dipole forces are stronger than even hydrogen bonding now that's not an absolute statement for some of the strongest hydron bonding and some of the weaker ion dipole forces it actually flips but in general ion dipole forces are usually stronger than hydrogen bonding so if we look at the chloride ion now so chloride being negative is actually going to want to interact with the hydrogen's in the water molecules so with the hydrants being partially positive : it's just some variable number of water molecules in either case so but this is what ion dipole interactions are and it always implies you have a mixture so which is why we didn't bring it up before cool so sometimes people talk about dipole induced dipole and things of a sort I'm just gonna leave that out some people just lump those together into London dispersion forces so also just wanna really quickly talk about what people call van der Waals forces so van der Waals forces depending on who you talk to means something a little bit different so now van der Waals have been customarily become dissociated with London dispersion forces and most the time people when they talk about van der Waals forces they mean a lot of dispersion forces but technically van der Waals forces encompass all of the intermolecular forces however more people just associate with limit aspersion forces and that's probably what you should do as well so odds are that's what your instructors going to use and that's what your TAS are using and that's what your friends are going to be using and so things are sort so you probably should just think that van der Waals forces is synonymous with London dispersion FYI so I'm do a couple more comparisons on some intermolecular forces before we move on there's some comparisons again just a little bit tricky here and stuff like that so we've seen that kind of hydrogen bonding rules the day most of the time for pure liquids again ion dipole force is only being present in mixtures we want to compare some other pure liquids and see what's going on here as well so if I give you again say ch3oh ch4 and h2 we already know that strongest intermolecular force was ch3oh because he's got hydrogen bonding and the others don't and so hydrogen bonding is usually king when you're comparing pure substances cool what if you don't have hydrogen bonding though so let's just say you don't and you've got something like this and you might be like well Chad that might have an fh bond well actually if you look at the structure you've got carbon bonded to the three hydrogens and carved onto the fluorine but the hydrogens none of them are actually bonded to fluorine in this truck there's no hydrogen bonding but there are dipole-dipole forces and align and dispersion now if you compare these three so you've got one that's polar the other two aren't but the truth is if you don't have hydrogen bonding polarity is probably not the most important thing unless all the molecules are the same size if they're the same size then they're gonna have roughly similar London dispersion forces and in that case dipole dipole could be the distinguishing factor however if they're not similar in size usually the biggest ones just gonna have enough London dispersion to even overcome any lack of dipole-dipole forces well in this comparison that's not a problem because ch3 F is the biggest molecule and the only polar molecule and he was gonna win no matter what but let's say we change this comparison up a little bit so instead of h2 at the bottom here let's say we go ch3 ch2 ch2 ch3 and now we have a problem because this last molecule with only CH bonds those are nonpolar bonds and if all your bonds are nonpolar your nonpolar so non polar nonpolar but the top ones polar and a lot of students think that the polar one should have the greatest intermolecular forces because dipole-dipole forces are stronger than lemon aspersion well they're stronger if you're two molecules of the same size having those extra dipole-dipole forces or the greater dipole-dipole forces would be the difference but if your molecules aren't that similar in size so in this case if this molecule is probably you know at least 20% bigger than this guy and he's way bigger than that so that's gonna be significant enough to give him enough extra London dispersion forces to overcome the fact that the CH 3f also has dipole-dipole forces and so if I asked you between these three compounds which one's got the greatest intermolecular forces you're supposed to realize that one hydrogen bonding is off the table none of them have it and as long as that's the case size is more important than polarity if you have any more than like a 15 20 percent difference in size size is more important than polarity the only time polarity will become the determining factor is if all your molecules are nearly identical in size and so in this case when I look at these are most likable he's the biggest by a fair amount he wins he's gonna have enough London dispersion forces to even make up for the fact that he doesn't have dipole-dipole forces and CH 3f does so greatest intermolecular forces is his last ch3 ch2 ch2 ch3 and because he's got the greatest intermolecular forces he'd have the highest melting point the highest boiling point the highest surface tension the highest viscosity but the lowest vapor pressure coolest one last little tricky comparison here and I chose two hydrocarbons specifically because they're both nonpolar and if they're both nonpolar what's the only intermolecular force they have hopefully you said London dispersion forces and that is totally true and so we want to compare their size and surface area and usually it comes down to size however these guys are isomers they both have the formula c5h12 and they both have a molecular weight of 72 they're identical but it's going to come down to surface area then and so in this case one of these has a much more compact structure than the other so one of them is kind of like when we crumple that piece of paper up so and it gets a much more compact structure and that's this branch structure right here branching kind of compacts the structure making it have less surface area and less surface area here is going to mean less lunna dispersion forces and so the top molecule is gonna have the greater lemmon dispersion forces if you think about like your intestines your intestines are a long skinny tube with lots of surface area and then even on inside there you got the villi that kind of have tons of surface area being long drawn-out structures and stuff like that so but once you get a lot of branching that really compacts the structure making it have less surface area so if I asked you which of these two molecules has the highest boiling point melting point surface area I'm sorry surface tension or viscosity it's definitely the top one the one with no branching and it's in its formula here so but again this usually comes into play with isomers that have exactly the same molecular weight so if one of these was a little bit bigger than the other I probably just wouldn't make it about surface area I just make it about size cool these are in molecular forces these are some of your common comparisons and you probably want to run through several examples to make sure you really got this I tried to make sure I gave you an example of every kind of crazy comparison now one thing we didn't talk about was comparing something nonpolar that's really big - something with hydron bonding and the reason we didn't is that's a tough comparison so to be big enough to have enough London forces to overcome a lack of polarity maybe 20 percent bigger than the polar molecule and you'll do it but to have enough you know for a non-polar mall you'll be big enough to overcome the fact that you know once have more animal forces than something with hydrogen-bonding it might have to be you know three four or five times bigger not 20% bigger like three hundred four hundred five hundred percent bigger so to speak and so but there's no like strict cutoff and so if they're gonna ask a question like this they're gonna make it huge they're gonna put something instead of like you know here like ch3 ch2 ch2 ch3 it's not a five carbon chain they might put like a 20 carbon chain or a thirty carbon chain and if you compare that to water well it would win even though it's nonpolar but it has to be significantly bigger and so and how significant is that there's no strict guideline for cutoff and so usually if they're going to give that comparison they're just gonna go overboard with it and make it ginormous so if turnout if you got like a thirty carbon chain that's like a wax and a wax is a solid at room temp even though it's nonpolar so whereas waters a liquid at room temp even though it's got hydrogen bonding so those are the least likely comparison you're gonna see usually hydrogen bonding is just gonna win the day all right next on our list here we're gonna take a little time to look at phase diagrams and we'll look at a standard phase diagram and then we'll see what makes water and carbon dioxide unique a little different from the standard so this case we do a plot of pressure on the y axis temperature on the x axis and it separates the three phases and we've got solids up over here liquids here gases here so and you supposed to definitely know what regions those lie in and kind of make sense if you kind of pick a point here and it's basically I'm gonna pick the point of one atmosphere and if you go from low to high temperature will increase along that x-axis here you'll see what that will go from solid to liquid to gas as we normally experience with most substances like water as you increase the temperature you go from solid to liquid to gas so it kind of makes sense where they lie and stuff like that in that regard cool you should be able to identify all your different phase changes then as well so if I tell you you've got a solid right here you let's say I've got a substance right here at that particular pressure and temperature you'd know it was a solid but if it crossed the line right here and I drew an arrow and said hey identify the phase change you should know all six of your phase changes and what they would look like on this kind of a diagram that arrow right there from solid to liquid we call melting or fusion the reverse process would be called freezing or crystallization if I went from liquid to gas that would be boiling or vaporization and from gas to liquid that would be condensation solid to gas sublimation gas back to solid either deposition or more specifically vapor deposition so know how to identify all your different phase changes on a graph like this so you should also know something about these lines these are called the lines of equilibrium because you have two phases in equilibrium with each other when you're on these lines so if I told you that we had some water and this water was at negative 10 degrees Celsius you'd know that was below the freezing point and 1 atmosphere and negative 10 degree Celsius pacifically by the way and you'd know it was ice solid but if I told you that you had some liquid water at one atmosphere and positive 10 degrees Celsius you'd know oh now I'm above the melting point of freezing point and it's going to be in the liquid phase but if I told you that you had some water at 1 atmosphere right at 0 degrees Celsius well then you'd know you're at the phase change temperature and you'd be right on this line of equilibrium and when you're on that line of equilibrium when you're right at the phase change temperature you have the two phases in equilibrium and so in this case right at zero you have solid and liquid you have liquid water and ice in equilibrium together cool so that's why they call these the lines of equilibrium so anywhere along this line over here you have liquid and gas and equal living together and anywhere along here you have solid and gas in equilibrant again and then finally right where they all meet right here so you have all three phases in equilibrium so that gets a special name and we call it the triple point and aptly named because all three phases are in equilibrium at only that one point on the entire phase diagram all right a couple of the things we should identify and we kind of already started here at one atmosphere we can't just drag this line across so you can see the melting temperature or freezing temperature right there at one atmosphere and then the boiling temperature and I guess you could call it the condensation temperature but nobody really says that so right there as well at one atmosphere and as long as you're talking about them at one atmosphere we call those the normal points and so this one right here that's your normal melting point and this one right here that's your normal boiling points cool so you should know that when we say normal melting point or boiling point we mean the melting a boiling point at one atmosphere and it can identify I'm on a phase diagram like this as long as the y axis here is labeled to show you where one atmosphere lies all right one more important point here and if you look at the solid liquid line of equilibrium it just goes to the top of the graph there's no cut off here just goes till however much graph you want to show it just keeps going but that's not true about the liquid gas line of equilibrium that liquid gas line of equilibrium stops right there and we call that the critical point and that critical point has a critical temperature and a critical pressure associated with it or right at that point and that critical point it's kind of an interesting point and a lot of students gets lost in them what that actually means and but the idea is that beyond that critical point there is no liquid gas phase transition so and that's kind of weird and it's kind of hard to visualize what what does that mean and stuff like that so let's kind of look at this for a second so let's say we've got a gas right here gas right here so and let's say we take that gas and we jack up the pressure well if we move up in pressure at some point we're gonna cross that liquid gas line of equilibrium and it's gonna turn into a liquid and so what's happening is you Jack the pressure up you're forcing the molecules closer and closer together and if they get close enough so the closer those molecules are the more they feel they're attractive forces if this is water we're talking about the hydrogen bonding they feel so and at some point if they get close enough the attractive force is just gonna win and all the sudden all the molecules are just gonna condense into a liquid so and it's a very sudden thing it's not a very gradual thing you know as you jack up the pressure if the gas gets smaller smaller smaller and that's gradual but once you hit that line of equilibrium rapid reduction in volume when condensation takes place and its really really visible you should be a professor who do this lovely demo and he'd take one of those big 55-gallon drums and before class started and before any students had come to lecture he would fill it full of super heat steam and then he'd sealed the top and then he would take that 55-gallon drum and just wheel it off and throw it in the corner of the room so and then students would start filing in and he'd start lecturing so in whole time that 55-gallon drum is just getting its temperature lower and lower so for that superheated steam and when it hits a hundred degrees Celsius and crosses that line of equilibrium as the temperature gets lower and lower and lower all of a sudden and it collapses into a liquid and that whole 55-gallon drum implodes all of a sudden now there's still pressure from all the air pushing on the outside but there's no more steam or leaves very little steam on the inside pushing back and the whole thing implodes makes a ginormously loud sound and makes a great demo scares the crap out of everybody it's it's beautiful so but it serves a purpose as well so if you look at what's happening now if we lower the temperature well if we lower the temperature now the water molecules are moving slower and slower and slower and at some point they will no longer have enough kinetic energy to overcome the attractive forces and the attractive forces win and it collapses into a liquid yet again all right so if we look at this critical point now so what if we take this same gas we've had and instead of jacking up the pressure or lowering the temperature in crossing that line of equilibrium let's say we just increase the temperature and then we increase the pressure and then we decrease the temperature and then we decrease the pressure and at some point here we've definitely turned into a liquid but where was that point I mean as soon as I crossed a line of equilibrium then it collapsed into a liquid it was an obvious point especially like in that example with a 55-gallon drum the whole thing just imploded in an instant but here there was never an obvious point when we became a liquid no we did but when was that because as we jacked up when we increased the temperature here the molecules moving faster and faster and faster and as we increase the pressure they got closer and closer and closer and closer but no condensation taking place and then we lowered the temperature so then the molecules moving slower and slower and slower and then we lowered the pressure back down but the question is when did it become a liquid because there was no obvious point where's like oops there it is that's that's a liquid now because it never went through a phase transition beyond the critical point here so at any temperature bigger than this the molecules will have so much kinetic energy that it doesn't matter if you jack up the pressure and get them really close together they will have enough kinetic energy to overcome any of the intermolecular attractive forces in those molecules and it's never just gonna condense into a liquid in an instant and go through a phase change so on one thing you should know is that critical temperature then depends on the intermolecular forces and molecules that have more intermolecular forces tend to have higher critical temperatures as well so another way we can relate this back to our discussion on intermolecular forces here all right but that's what a critical temperature is the temperature beyond which there is no liquid gas phase transition and so if you're down here you got a gas and if you're over here you have liquid but if you're out over here you have what's called a super critical fluid supercritical meaning above the critical temperature fluid instead cool that is your critical points all right let's take a look and see how in the world water and carbon dioxide are just a little bit different here all right this is water so and I drew in blue what makes this so different I just should have drawn it in black but I want to make sure it was really obvious what was different here and we still have solid here liquid here gas here that's all in the normal places and so it's usually and I've even actually amplified the slope here it's usually very subtle here but the key is this the solid liquid line of equilibrium has a downhill negative slope so it's mathematically a negative slope where is your typical compound over here that line has a positive slope now I've amplified it here to make it so it's obviously negative but it's really subtle so but it is indeed a negative slope and that is unusual there's only couples you know substances I can think of and water is one of the two that does this almost every other compound does this okay so what's going on here well everything that is unusual about water usually comes down to hydrogen bonding and that's going to be the case here as well so let's say we've got a substance here so now let's say we've got that substance right under these conditions right here and we jack up the pressure okay well right now it's a liquid and if we jack up the pressure at some point we're gonna cross that solid liquid line of equilibrium and it's gonna turn into the solid and the idea is that you're putting pressure on this this liquid you're squeezing in on it so and generally as you put pressure on substance it wants to get more and more compact more and more dense in structure and for most compounds the solids more dense than the liquid and so it's going to cross that line and turn into a solid so however if you've got water as a liquid and you jack up the pressure it injustice a liquid because that's what's unusual about water is the liquid again is more dense than the solid when you freeze water it expands the solids less dense so but you can do something interesting on the other side instead is you're gonna have some solid right there and you can jack up the pressure and all of a sudden you cross that solid liquid line of equilibrium and it turns into a liquid instead and this is really unusual so this is really important this is just a reflection again that ice actually is less dense than liquid water and this is super important now one it has to do with you know and like the oceans freeze and stuff like that it floats to the surface and it insulates the rest of the ocean from cold weather and prevents the oceans from freezing through and through which would kill all life and you know it's kind of important I guess killing all life and yadda yadda but the really important thing is that this is vital for the game of hockey so if you look at the way hockey works we like to skate on ice so and it turns out that wet ice has far less friction than dry ice think about a slip and slide and taking a run on a slip and slide and jumping out to onto that slip and slide and then realizing that you forgot to turn your hose on you kind of get a pink belly like no other right so because water on the slip and slide decreases the friction and you slide much better that the same thing is true on ice when ice is a little bit wet it's much more slippery ie it just has less friction and stuff like that and so the same thing is true in hockey and what happens in hockey recall that pressure is force per unit area is that the hockey players weight pushes a force down onto the ice and it does it right on the blade of his skate that's why the blades of those skates are really really small they have a small surface area which leads to a large pressure right under the skate and as a result the ice right underneath the skate actually crosses that line of equilibrium and melts leading to less friction and you can skate a lot faster I always thought it'd be kind of a cool prank to play on an NHL team or something like that and so sneak into their rink and Mel to all the rice and get rid of it and replace it with dry ice so in that way dry ice doesn't have this negative slope it has a positive slope so you're placed with dry ice and they're gonna skate around and they're like why do I feel so slow today so and if you you know really funny and a good laughing you know until some of that dry ice I guess sublimed and they all have fixated and died but up until that point it would be really funny I'm telling you all right so that's water that's what makes water unusual and the last one that's unusual is dry ice it's carbon dioxide and again you notice with dry ice here carbon dioxide let's write that in that solid liquid line of equilibrium does indeed have a positive slope that's not the unusual part here what's going to be unusual is where we find one atmosphere so for one atmosphere we're below the triple point most compounds one atmosphere lies above the triple point and so most compounds as you heat them up they go from solid to liquid to gas but not carbon dioxide for carbon dioxide you heat up solid dry ice and you never get a liquid at one atmosphere it goes straight to the gas and sublimes this is why they call it dry ice right it never feels wet because it never is it's never liquid if you want to have liquid carbon dioxide it doesn't exist at one atmosphere it only exists at elevated pressures so this is the unusual part of carbon dioxide so one atmosphere lies below that triple point so if I asked you for the normal melting point of water you'd say zero degrees Celsius now if I asked you for the normal boiling point of water you'd say 100 degrees Celsius and again normal meaning the melting point and boiling point at one atmosphere but if I asked you for the normal melting point and boiling point of carbon dioxide it's a trick question because it doesn't have one there is no melting point or boiling point at one atmosphere for carbon dioxide and therefore it doesn't have a normal one in this case it has a sublimation point instead cool those are your phase diagrams no the standard no it makes water and carbon dioxide unique as well and once again if you are looking for some good practice problems on this stuff check out Chad's prep comm and my ultimate general chemistry prep course and if you learn something in this video like it share it to end scribe you know the drill hope you're having a great day