Transcript for:
Understanding Intermolecular Forces and Phases

okay so for today's lecture we're going to talk about chapter 11 which focuses mostly on intermolecular forces between different substances as well as the different phases that substances can take depending on the set of circumstances that they are in an intermolecular force is an attractive force between two separate molecules okay it's not a chemical bond but it's an attraction so if we think of something like water is a very good example of this in you know if we had if we had water in a situation where it wasn't affected by something like gravity we could see its intermolecular forces at play perfectly okay so on the space station we have you know a first-hand view of this or maybe I guess a second hand because we're not there ourselves but we could see what happens if you spill something like water in a zero-gravity situation what happens is that it forms these little spheres okay and that is because of the intermolecular forces between the different water molecules intermolecular forces exist in every particle that makes up matter every single particle has an intermolecular force with other particles around it and the different phases were the different states that substances exist them depend on the strength and type of the intermolecular forces that they experience okay so I'll make sure look at the slide here because this is a weird thing to look at usually when we have materials gases would be the least dense material liquid would be second and then the most dense material would be a solid version of it right if we were going to move through the phases of a particular substance water is an anomaly here we can see that its density in the gas phase is very very low as we would expect if we're comparing the liquid and the solid densities here are the density of liquid water is greater than the density of solid water which is ice right and that's why ice can float on top of water that doesn't normally happen with a solid and a liquid and that has to do with the intermolecular forces that are at play which we can get into more specifically in this chapter okay here are our three states of matter and their corresponding intermolecular force strength okay so if we have gases which have a low density they have very weak intermolecular forces that means that they can separate from one another easily that is why they can take up the shape and the volume of the container that they're in because there's nothing holding them together liquids on the other hand which begin to have a higher density than gas certainly higher than gases like that's a definite they're intermolecular forces are in the middle of these in general okay so they have moderately strong intermolecular forces they can still take the shape of their container they can separate enough from one another to be able to flow that's what that's why they are able to take a shape of the container that they're in but their volume is a fixed volume and then we have solids which have very strong intermolecular forces the intermolecular forces are so strong that the particles can't separate for another and they have a definite volume and a fixed shape okay if we see this visualization here of a liquid we can see that the molecules are packed in pretty tightly against one another even though that's the case they still have the ability to move around past one another so they they do they are close together where they can still move around the fact that they can move around is what gives them the ability to flow and that is why they are able to take the shape of this container they're not in a fixed shape but the fact that they are so close together makes them incompressible incompressible so if we try to put this amount of water in a smaller container it wouldn't fit we couldn't squish those molecules in any closer together than they are gases on the other hand are both able to move freely past one another and at great speed and they're constantly in motion but they also have space between them so you see all these gaps between the gas particles that means that gases are compressible and we can sort of squish them down into smaller spaces and they still exist just fine like that okay they are compressible and because the particles are not in close contact they can move freely and they that's what gives them the ability to fill the volume of the container that they are in this is an example of what we mean by compressibility just so we can visualize it better if we have a liquid sample in a space and we try to push down on this on the the ceiling of the container try to push down we couldn't push it down any lower right there's no there's no big air spaces or big pockets of empty space between these molecules the volume is the volume the gas sample on the other hand we can push down the top of this and the gas sample still fits easily inside the new volume it just has less empty space between the particles now solids on the other hand are in movable okay they have got a fixed position the particles of a solid are not going to move past one another but they still might vibrate in place because every atom even atoms that are part of solid matter they are still in constant motion in this case a motion is just a vibrating motion and we have two types of solids we have crystalline solids and amorphous solids are crystalline solids are nice ordered solids they have a repeating pattern that exists within them the amorphous solids have no kind of pattern all right they're just kind of thrown together however but they are still fixed in position okay now the close packing here in both cases makes solids obviously and compressible but it also means that they are because atoms are fixed in position they are rigid so they have their own fixed volume they don't take the volume of the container that they're in and solids don't flow right they stick together examples of crystalline solids or or our solids that are very orderly and patterned are things like salt and diamonds okay this is where you would see that repeating order regardless of the point in the sample that you're looking at examples of amorphous solids are things like plastic and glass and these again are things that don't have any particular pattern to them so just a quick concept check what state of matter is compressible solid liquid or gas so hopefully we are all on the same page here so now we can talk about these changes we are going to focus on phase changes a lot during this chapter and the phase change is anytime a material goes from one of our phases into another phase so I mean that makes sense right so a solid would melt once heated a liquid would boil when it's heated those are phase changes and these transitions can be induced by temperature changes or by pressure changes okay so temperature and pressure are the two ways that we can force a phase change from one state to another state here's another conceptual connection concept check which of the diagrams below best depicts the vapor emitted from a pot of boiling water okay so this is liquid water here and which one is going to be boiling water we would expect that if we were forcing a phase change from the liquid state to the gas state that our water molecules would not separate from themselves they would still stay as water molecules but if they're forming gas from a liquid they should be more spaced apart there should be a lot more room between them okay so choice a here is the only one that still has our water molecules together as they should be and now they are much more spaced out so that is a visual depiction of boiling water now intermolecular forces are what determines how much how much energy it's going to take to separate molecules from one another because they they are the attractive forces that keep things together okay some of the particles are attracted to each other by electrostatic forces right most of the particles these are the positive and negative charges that happen within the molecules the intermolecular forces are going to actually determine the state the substances in because if you think back to that chart that we already looked at gases have very weakened and molecular forces liquids have moderate ones and solids have strong intermolecular forces and the strength of the attractive forces is going to vary depending on the material we're looking at right the stronger the attractive force or the stronger the intermolecular force the more those particles are gonna resist moving okay so we already talked about that okay Mary talked about some of this stuff so let's talk about why molecules become attracted to one another this is gonna have to do like we said with these electrostatic charges okay so we have molecules and within the molecules themselves there can sometimes be these positive and negative charges all right so in this case we have a polar molecule this is water and water is polar so we know that it has a partial negative charge on the oxygen and a partial positive charge on the water on the hydrogen rather so the partial positive hydrogen becomes attracted to the partial negative oxygen here it doesn't form a bond or anything like that but there's an attraction between them and they kind of stick together okay like like a magnet sticks together no it's easy to see this when you have a polar molecule because you can understand that you have this positive and negative end and they might be attracted to one another but it turns out that even nonpolar molecules can have temporary charges or temporary dipoles and we'll look at that in a minute but what is true also is that the larger charge that you have the stronger the attraction is going to be between the two molecules and the further the distance between two things the weaker the attraction is going to be okay now if we are trying to compare these attractive forces to bonding forces there's no comparison because bonding forces real real chemical bonds are just so much stronger than intermolecular forces and but you know they still exist though and they still affect the behavior of the material that we're looking at okay so these are our three main there's there's four total forces that we're gonna look at these are our three main ones though okay because these are forces that can exist within pure substances the fourth one that we're going to look at is still an intermolecular force but we need like some kind of mixture for that to happen so our first is called dispersion forces dispersion forces happen between any material any any matter that exists has dispersion forces in it they don't require a permanent dipole okay dipole-dipole forces do require polarity in the molecules in order to exist and then we have hydrogen bonding which is basically a an incredibly strong dipole-dipole bond that only exists when hydrogen is attached to a very electronegative atom so now we're going to go through these one by one so dispersion forces work because of what are called temporary dipoles that form within nonpolar molecules okay you can have these areas within a nonpolar atom or a non-polar molecule that sort of very very temporarily have an excess area of electron density and a depleted area of electron density okay and this is because if we look at this picture here um electrons are in constant motion okay so we have these two helium atoms right here well in helium atom 2 over here its electrons both ended up kind of next to each other okay so this is pretty negative now when that happens the helium atom that's next to helium 2 well its electrons are not going to want to come anywhere near this negative charge here okay so they kind of back away and what happens then is that the nucleus of the helium becomes completely exposed to the electron force the electronegativity of helium two and these two things become attracted to one another okay so we have a completely non polar atom which ends up forming basically an attractive force between the proton of one and the electrons of another this is called an induced dipole because it doesn't exist without extra help alright and that help is the pooling of these electrons in the helium atom over here which basically push the other electrons away and leave the nucleus of our of this atom exposed to the negative force of the other electrons okay so these attractive forces are caused by temporary dipoles and we call them dispersion forces or London forces and every molecule and atom has them okay as these temporary dipoles occur they continue to induce dipoles and surrounding molecules right because now as this dipole this temporary dipole comes into place well now these electrons are going to kind of migrate towards one another and now they'll form an induced dipole with a helium atom that's next to them okay it's a temporary uneven distribution of electrons and it causes a temporary attraction it's not a permanent partial charge okay and in another few minutes they might reposition themselves and now these electrons might be over here forming a temporary dipole with a helium that's up here okay but they are but the attractive force is still present between the different nonpolar material we can imagine that the size of these temporary dipoles are going to depend on things like the volume of the electron cloud and the molar mass of the material that we're looking at because both of these factors have to do with the amount of electrons that are present and the amount of electrons I'm going to switch back again that can participate in this forced dipole say we had a material that had kind of ten electrons bunched over here okay if I attend electrons bunched over here and it was pulling on this nucleus well that's gonna be a stronger force than just two electrons pulling on this nucleus okay so the the size of our molecule is going to take it's going to have a big effect on the on the pole between the nucleus and the electrons it's a little larger the molar mass of the material that we're looking at those stronger those attractions are going to be in the shape of the molecule is going to matter as well because the more surface contact we have between the two particles the bigger the induced dipole can be because you'll have more areas where the electrons on one can be attracted to the nucleus of the other material all right so I'm going to show you a picture of that in a minute I'll come back to this okay I want to show you this one so we have this both of these are examples of pentane but this is a straight chain pentane and this is a branched pentane so if we have these straight chain isomers these straight chain molecules if you can imagine at every point here if we like laid one of those molecules on top of another molecule we could have like one two three four five six contact points between them okay so that's six contact points where we could have the electrons from one of them be attracted to the nucleus of the other one right so that increased number of contact points means that you're going to have stronger intermolecular forces in this case where we have this branched molecule it almost looks like it's trying to make a sphere or something like that but we're gonna have significantly few areas where they would be contacting where they would be able to be in contact at the same time alright so the fact that there are fewer contact points here means says we're gonna have fewer electrons to nucleus attractions which means that we're going to have weaker dispersion forces and we see that reflected and the boiling points of them okay so the boiling point of the neopentane is 36.1 degrees Celsius the more abnormal pentane sorry the straight chain one the boiling point of the neopentane which is the branch chain that only has a boiling point of nine and a half degrees Celsius okay because there's so many fewer points of contact I now let me go back to where I left off before I got off on my tangent this is an example of how the size of the different atoms affects the boiling point so we can see if we compare the boiling point of helium which is four point two Kelvin to xenon which is 165 Kelvin as we increase in molar mass we're increasing in the of our dispersion forces and that increase in strength of dispersion forces is reflected in an increase in boiling point and this is because we have more electrons available in the bigger molecules so in something like xenon we have significantly more electrons that are available to pull on the nucleus of corresponding xenon atoms that makes the dispersion forces stronger and that makes its boiling point higher so the stronger the attractive forces between the molecules the higher the boiling point will be that's a good rule to remember and this is just a visualization of the boiling points of different alkanes these are all straight chain alkanes if you see but if we compare the 5 carbon boiling point to the 9 carbon boiling point you see how much bigger it gets and that's because we have more electrons present to participate in the dispersion forces we just already talked about this so the straight chains are going to have more surface to surface contact which will give them stronger dispersion forces so according to our conversation that we just had which halogen would have the highest boiling point chlorine bromine or iodine iodine is the biggest it's the most massive it has the most electrons it'll have the highest boiling point all right so now we'll move on to the dipole-dipole forces they exist between molecules that have permanent dipoles in them already okay so if you look at a molecule and that's why you spent so much time looking at this in the previous chapters if you have a polar molecule then you have a molecule that will exhibit dipole-dipole forces okay they the dipoles in them will always be present okay it's not like the temporary induced dipole that we looked at in dispersion forces these are permanent okay um and if you look at well you can look at any example here but I guess we can imagine it maybe better on this electrostatic map this red area shows a negative charge of that oxygen that's present in this molecule here and the more blue areas are the positive charges so we have our negative charge focused up here so if we had many of these molecules that existed next to each other what would end up happening was the positive ends of these molecules would align themselves so that they were nearly in contact with these negative ends and there would be an attraction between the negative end and the positive ends and those are dipole-dipole forces okay so this is a visualization of what I was just talking about so the negative end of it would be attracted to the the most positive end and this is sort of how they would align themselves in a container together or something like that okay um again these are not a chemical bond but they are still I really like to think about intermolecular forces in terms of magnets and you know how magnets come in different strengths and you can have like a really like a weak magnet that won't even stay on the fridge and then you Crump a super strong magnet that's how you can think about these different intermolecular forces so like dispersion forces they are like magnets but they're very weak magnets and then when you get to dipole-dipole forces those are a little bit stronger like those are gonna hold up like a kids school or on the fridge and then you have you know one step past this is our hydrogen bonding and that's gonna be like super strong magnets so it goes kind of in that in that trend now when you have you know stronger attractions like we've already talked about it's going to take more energy to separate them so this is going to be true or your dipole-dipole forces as well the higher the boiling point of any liquid the stronger the intermolecular attractive forces are going to be I know that I'm saying this over and over again but I really want that trend to be like hammered into your head by the end of this we have to remember though that when we are breaking apart intermolecular forces we are not breaking the covalent bonds that make up the molecules themselves the molecules are still going to exist as the molecules they're just going to exist in a different state if we're doing something like boiling water okay so we can look at the dipole-dipole effect on different molecules we have two hydrocarbons they are basically the same molar mass okay so if we just use molar mass to compare these and we didn't know anything about the structure of them or the pool area but we would look at f/8 and be like oh it's a little bit heavier who knows maybe it'll have a higher boiling point at the very least they should be the same okay but then we look at the structure and we're like oh we have this oxygen here that's gonna be pretty electronegative oh and so this molecule is obviously going to be a polar molecule so the dipole-dipole forces that are present in the formaldehyde raise the boiling point so it's negative nineteen point five degrees Celsius ethane is boiling point is negative eighty eight degrees Celsius that's significantly lower and that's all attributed to dipole-dipole forces because they have very similar dispersion forces solubility is also going to depend on intermolecular forces because similar material is going to dissolve similar material like like dissolves like um things that have polar groups are going to dissolve polar groups things that are nonpolar are going to dissolve nonpolar things examples of what are called hydrophilic groups meaning that they are polar or they would be okay with dissolving in water okay these are groups like OHCHR - yo you notice the trend here is that they're all going to have these electronegative parts that are going to make them polar on the other hand the hydrophobic groups which would like to avoid things that are associated with water which is polar these are your nonpolar you know these are nonpolar bonds here your carbon hydrogen bond your carbon carbon bonds those are nonpolar bonds they don't want anything to do with polar water molecule so when you have molecules that have a hydrophilic and a hydrophobic part when you're trying to dissolve them in water it comes down to the like the level of the intermolecular forces between the polar groups and the water and the nonpolar groups on the other molecule itself so we have to take both of those into account we're figuring out if something is going to dissolve in water we can look here at an example of us trying to dissolve pentane and water if we do that the pentane has no polar groups to it at all it's a completely non-polar molecule and because that's true it has no reason to want to form any kind of intermolecular forces or any kind of attraction with these polar water molecules so it's happy what happens is they end up layering one on top of another the water is more dense so it sinks to the bottom and the pentane just sits on top of the water and it actually forms a layer that you can visually see liquids that form these layers and that do not mix together are called in miscible liquids they're liquids at do not form a modulus mixture when you add them together well let's look at this electrostatic map of a pseudo nitrile a pseudo nitrile is polar we can tell that from the map our negative end is over here our positive end is over here from the map determine how they would interact with each other in space okay like which of these orientations look like they would be most correct now remember when you're trying to do this you would want the negative end of the molecule to be very close to the positive end of the next molecule that's what we're looking for okay so in this case structure the structures in a most closely represent that our next type of intermolecular force that we're gonna look at is hydrogen bonding this is our very strong magnet for hydrogen bonding to occur we need our hydrogen to be attached to one of these three molecules okay because these three are that are just very very electronegative so we have oxygen nitrogen and fluorine with hydrogen's bonded to them but the hydrogen bonding does not occur within this molecule itself it occurs between other molecules that are around this one okay and the reason why hydrogen bonding is so strong is because hydrogen has no other electron right it's just it has the one single electron in it so when it attaches to an electronegative atom its electron gets pulled towards the electronegative atom and then what happens its nucleus is completely exposed or die's shielded so that means it has a full positive charge that's just sitting there out in the open what that's gonna do is that's going to pull the you know the negative charges that are around it towards it that forms a strong attraction okay we can look at hydrogen bonding that happens within water now the hydrogen bond again it's not between the oxygen and the hydrogen that is a covalent bond okay but the oxygen at this point has basically pulled both of the hydrogen electrons towards itself so the oxygen is feeling very negative over here and the hydrogen's that are in the corresponding water molecules have exposed protons so they're very positive so the positive charge of this hydrogen and the negative card to this oxygen is very strong attraction between them and that ends up being that strong magnetic force that I was talking about okay now water is not the only thing that experiences hydrogen bonding we can see how should an bonding happening and something like ethanol as well it has it's an alcohol group so it has an O H over here the hydrogen from the O H would be hydrogen bonded to the oxygen from a different ethical that different o H group and again this is not a chemical bond it's just an attraction a very very strong attraction they're stronger than dipole-dipole forces they're stronger than dispersion forces they're the strongest intermolecular force that happens with impure substances okay that being said they are not as strong as chemical bonds covalent bonds are dramatically stronger than hydrogen bonds okay you see this two to five percent the strength of covalent bonds but you know still considering the fact that the water molecules are not chemically bonded together the hydrogen bonds still are you know they're strong enough to affect the properties of the materials that they existed so if we compare these two things we have ethanol and dimethyl ether have exactly the same older mass okay we have I mean if you're looking here at the structure of them we can see the oxygen in the middle here this oxygen is going to be bonded to a carbon on this side and a carbon on the side so that's not going to be hydrogen bonding here the hydrogen bonding is in this Oh H group for the ethanol look at the difference between boiling points your dimethyl ether is negative 22 degrees Celsius just by having the hydrogen bonding there the boiling point of the ethanol increases - seventy eight point three degree Celsius that is an incredibly high jump right it's nearly a hundred degree difference between them so these materials that are listed here they all hydrogen bonds so you would expect their boiling points would be higher than they would be otherwise even if we look back I'm gonna flip back to the other slide we have ethanol over here right it has a boiling point of seventy eight point three degrees Celsius it has a molar mass of forty six grams per mole water only has a boiling point of or I mean water only has a molar mass rather of 18 grams per mole its boiling point is still higher than the boiling point of ethanol and that's because of the dramatic hydrogen bonding that happens within the water molecules because those bonds are so polar and because the hydrogen bonds are therefore much stronger than they are in other material so if we're gonna compare that to like a boiling point of nonpolar things right what do we have for our nonpolar molecules like our group for molecules we only have boiling points that are based on dispersion forces in Group four so yeah they increase down the column the dispersion forces increased down the column as our as our atoms get bigger and bigger and the boiling points increased down the column but they're still not going to be boiling points that compare to things that have dipole-dipole forces our polar molecules are going to have both dispersion forces because everything has dispersion forces and they're going to have dipole-dipole forces so they're going to have higher boiling points than any corresponding group for molecule and their boiling points are also going to increase as you go down the column now if we have a mixture we can have one more type of intermolecular force that comes into play okay but this this has to be you're mixing two different things together you have ions from an ionic compound and they become attracted to the dipole of polar molecules okay these ion dipole attractions are one thing that determines the solubility of ionic compounds in water okay because they're actually able to you see they act like if you look at the NaCl down here the ion dipole attraction that's happening here is going to affect the chlorine from reforming a solid with the sodium okay because the chlorine molecule being very negative is going to be very attracted to the positive hydrogen atoms and it's going to surround itself with these positive hydrogen's right the sodium on the other hand is going to surround itself with the negative oxygens so now that it's done that it can't reform a solid with its counter ion okay so these attractions here excuse me these attractions here again are not chemical bonds they are just very very strong attractions okay these attractions are even stronger than hydrogen bonds but you need a mixture in order to have them present okay ah so let's summarize these intermolecular forces our dispersion forces are the weakest of them they're present and everything and the dispersion force is going to increase with increasing molar mass polar molecules in addition to dispersion forces are going to have dipole-dipole attractive forces okay you have to have a polar molecule in order to have a dipole dipole okay hydrogen bonds are a type of dipole-dipole interaction but they are the strongest of the dipole-dipole interactions and then the strongest type of intermolecular force that a pure substance can have to how the hydrogen bond your hydrogen has to be connected to a fluorine oxygen or nitrogen atom then you have your ion dipole attractions those are going to be present in any mixture of ionic compounds and polar molecules and these are even stronger than hydrogen bonds and they are going to be especially important when we are looking at aqueous solutions of ionic compounds because that is the that's the most fertile ground for having free ions available to pull those water molecules towards themselves okay so this is just a summary chart of the types of intermolecular forces and how strong they are so out of these substances which would have the highest boiling point we have ch3oh co and n2 okay over here okay start with n2 we see that it is nonpolar so it would have just dispersion forces Co is a polar molecule it would have dipole-dipole forces and dispersion forces and then ch3oh and has this alcohol group attached to it that alcohol group would undergo hydrogen bonding so it's got hydrogen bonding and dipole-dipole forces and dispersion forces so that's going to have the highest boiling points okay so let's look at this example here which molecules out of the three that are listed have dipole-dipole forces okay sometimes it's not that clear to see when you're just looking at the molecular formula like this so it's a good idea to write out a lewis structure for these and see if you can find you find from the lewis structure whether they're polar or not so let's go molecule by molecule will start with co2 if we draw out this structure Oh double bond C double bond o we know that the oxygen carbon bond is a polar bond because it has a significant electronegativity difference right two and a half to three and a half but the geometry is linear right so if we were looking at the vectors of the dipoles present in the carbon dioxide the vector would sum to zero so we would not have a net dipole so because the molecule itself is not polar it is not going to have any dipole-dipole forces present right it would only have this version for a ch2 CL we have if we draw the Lewis structure we see that it's a tetrahedral molecular geometry and that we have these two chlorines that are attached that are highly electronegative okay and we also see that when we look at these dipoles if we try to sum them together we would end up with a net dipole moment there because we have a net dipole moment that makes this molecule a polar molecule because the molecule is polar it's going to have dipole-dipole forces and if we compare this to ch4 we have a molecule that's tetrahedral its bonds are essentially nonpolar because there's only a small electronegativity difference between them and because it's nonpolar is not going to have any dipole dipole forces okay so hopefully that part makes sense we can look at some of the other properties that are affected by intermolecular forces now and we'll start with surface tension okay practically speaking surface tension is you know I'm trying to if you imagine like a pond or something like that and you see sometimes like debris that sits on the top of the water right like it's not necessarily floating it's just sitting on top of the water that's a visualization of surface tension okay water has very high surface tension this is actually a property that liquids have that results from their tendency to minimize their surface area or squeezed together okay so liquids try to do that by forming spheres so they try to bunch up together if you had that's why we talked about at the beginning if you have no gravity and you spilled water the water would sort of adhere to itself almost and it would form these spheres obviously if we have something like a giant pond or a giant lake water couldn't form a sphere but it does still try to get very close to itself okay um and this this is a good visualization of it too this is what happens we see that we things that are that are not very heavy right they can't break the surface of the water and that's because your water molecules are pulled so closely together to one another okay I like this description down here that the surface layer acts like an elastic skin so it almost allows things to float so paper clips are made of metal metal is gonna be denser than water but until the paper clip pierces the surface of the water it's going to sit on top of it okay and the reason that it can sit on top of it is because of the electro the electrostatic forces between the water molecules they have very very strong intermolecular forces they've got dispersion forces and dipole-dipole forces and hydrogen bonding all right so what do we got here the surface tension of a liquid is yeah that's a boring it's a boring definition of this one the surface tension is the energy required to increase the surface area to give an amount I think it's um you know I think it's it like makes more sense to think of it in the non mathematical terms like if you have you know a great deal of intermolecular forces you're going to have a higher surface tension than you would otherwise but I wait we do have to take into account is the fact that because you know you don't necessarily have um you're not necessarily being surrounded by molecules if you're at the surface of the container right because if you think if you're if you compare the molecules at the surface of a container with the molecules that are like in the middle of the container molecules that are in the middle of a container are gonna be 100% surrounded by other molecules um that's not true for our surface molecules right they're missing at least half of their surrounding molecules that top half where they're exposed to the air so they're going to be inherently less stable than your interior molecules and they are going to have a higher potential energy the stronger your intermolecular forces are the higher surface tension is going to be but if you raise the temperature of a liquid it's going to reduce your surface tension okay so things that are in warmer environments are going to have less surface tension this is because you are going to be speeding up the molecules themselves when you raise the temperature of anything you increase the kinetic energy of it increasing that kinetic energy means that you're going to speed up the molecules if the molecules are in constant motion that means that it's it's easier for them to not be next to each other okay and good surface tension is happening because you have these molecules that are kind of drawn very closely to one another that's how they form that quasi skin excuse me I'm sorry so what they're moving around then they're basically I don't know it would be like if you were stretching dough right and you stretch to get your stretchy or stretch and eventually it comes apart that'd be like the same same thing that's happening there viscosity is another property that is affected by intermolecular forces and if we have something that has a very strong set of intermolecular forces we have material that is very very viscous what is viscous it's the resistance of a liquid to flow so it's like very thick liquid very thick liquids or viscous liquids so if you're going to compare if you had a glass of water and you had a glass of like engine oil for a car and you poured those two out you would see very clearly that your water is much less viscous than your oil so that that's kind of what we're looking at there but we can also look at other liquids like if we were comparing our hydrocarbons again and we went from pentane to nonane which is again the five carbon to the nine carbon well we have point two four zero two point seven one one in terms of viscosity right the things that effect viscosity are intermolecular forces and the molecular shape okay so the more spherical the molecular shape of things are the lower the viscosity will be and that is because your molecules are going to roll more easily so it's a lot like what we were talking about with the dispersion forces when we're talking about straight chains versus branch chains straight chains have more surface to surface interaction which means they have more attractions spherical things or branched things have less surface to surface contact so they're not going to be as attracted and that's going to mean they don't stick together very much right viscosity it's almost like stickiness if we raise the temperature of things that is also going to reduce the viscosity that's going to increase again the speed of the molecules if you're increasing the speed of the molecules that means that the molecules have an easier time separating from one another and if molecules can separate from one another and overcome the intermolecular forces that are keeping them together that allows them to flow capillary action is another property that is affected by intermolecular forces capillary action is the actual ability of a liquid to flow up the like up a very very thin tube against gravity so if you put a capillary tube in a liquid the water would go up even though gravity says that water should go down and this happens because of two separate forces that are working together we have cohesive forces and these are the forces that hold the liquid molecules together and then there's also add adhesive forces and these are forces that are going to attract the outside of the liquid molecules like so any of the liquid molecules that are attached to the or that are in contact with the surface of the capillary tube there would be an attractive force between them we can actually look here visually at the different types of capillary action that depend on the diameter of micro capillary tubes that are being used you see that the narrower the tube is the higher the liquid is going to rise and that's because of those two forces that we were talking about okay if you have a thinner tube a tinier tube then more of your sample is going to be in contact with the outer walls of the tube and if more of your sample is in contact with the walls of the tube then more of your sample is experiencing the adhesive forces so that's going to allow it to rise further up the tube now in in this case here in this tube the bulk of your sample in the middle here is not experiencing adhesive forces so what ends up sitting here the smaller the tube gets the more of your sample is experiencing adhesive forces so the more that sample is going to rise up the tube um we're going to look here at the meniscus and now in general this is affected by your intermolecular forces we have two kinds of meniscus that we need to look at okay the meniscus is the curve and the liquid that's gonna happen whenever we have liquid that's in some kind of thin tube like a graduated cylinder or a test tube or a burette anything like that the meniscus forms because of a competition between the adhesive forces and the cohesive forces so it's sort of like depending on which one wins that's gonna affect the shape of your meniscus we're used to seeing a water meniscus okay because we work so often with aqueous solutions we're used to seeing this concave curve right this one here is a little bit more foreign to us because it's this convex curve and that doesn't normally happen in water solutions so we have this concave water meniscus because the adhesion to the glass is stronger than the cohesion for itself okay so instead of sticking to itself the water is going to start to creep up the sides of the glass tube because of the adhesive forces and that's what forms that proper u-shape and something like mercury in contrast the cohesive forces within the mercury are stronger than the adhesive forces all right so we end up with this like horseshoe shape right because the mercury would rather stick to itself then touch the sides of the glass tube and this is because metallic bonds so bonds that would exist within mercury metallic bonds are stronger than intermolecular forces so the metallic bonding that's happening within the mercury sample is stronger than all of the other intermolecular forces that are happening within the water so the mercury is going to stick to itself while the water is going to stick to the sides of the glass and form that concave shape okay now we're going to talk about some phase changes we're going to kind of skip into our phase change section phase changes have to do with molecular motion okay that's going to be the basis of all of our phase changes um and our molecular motion has a lot to do with the temperature that material is at and the pressure that the material is that okay because the those two factors are going to affect the amount of kind of kinetic energy that's available our first phase change that we're going to look at is vaporization and that is a transition from liquid into gas taking right now there's multiple things that are going to affect vaporization intermolecular forces are going to affect vaporization temperature is going to affect vaporization but also the surface area that's available is going to affect vaporization because the larger the surface area you have the faster your liquid will evaporate okay if you have a lot of high energy molecules and they are all existing at the surface of your sample at any moment they can have enough energy to overcome their attractive forces to one another and escape okay and every time that a molecule escapes and enters the gas phase so the more molecules that are present at the surface of a sample the more chance you have for them to escape the liquid state and become gas very percentage-wise only a small amount of your sample are going to have enough energy to escape but as you increase the temperature you're going to increase the available energy so that's going to increase the amount of molecules that are going to make this transition from the liquid state to the gas state condensation is a reverse process of vaporization and that happens when you have gas molecules that are losing energy because of molecular collisions that are happening okay as these molecules lose energy lose energy they actually end up reverting back into the liquid state they get captured by the liquid and then come back into the liquid state okay um sometimes instead of if you don't have like a liquid sample that they're getting trapped back into like when you notice on the surfaces you see condensation happening as droplets that's because these gas molecules are losing energy and then they start to stick together so then they start to form droplets of liquid on surfaces or around surfaces all right so they didn't quite hit the surface of the liquid but they hit the wall of the container near the surface of the liquid and then more low energy gas molecules started to hit the same area and they form these droplets so that that process is called condensation and it's the reverse of vaporization okay they're opposite processes if you have an open container your vapor molecules are going to basically exit and then keep going they're gonna exit and they're going to spread out and they're gonna spread out very quickly because that happens in general if you have an open container the net result is that your rate of vaporization is going to be greater than the rate of condensation so you're not going to be able to rican dense your liquid and maintain a constant liquid level okay in an open container there's going to be a net loss of liquid because as the liquid vaporizes it's going to sort of exit the area and not be able to reform into its liquid state if however you have a closed container and the vapor is not allowed to leave then the net results between the vaporization and the condensation is going to be equal and you will not have a loss a net loss of liquid whatever ends up vaporizing will eventually condense so that's kind of interesting when we talk about intermolecular forces and how they affect evaporation and condensation we have sort of I mean hopefully it makes sense because if you have very strong attractive force between molecules it will be hard to separate them from one another okay that means that it's going to be very hard for them to leave the liquid state and to go into the gas state so you'll end up with you know you don't end up with them with that much paper however if you have weak attractive forces then it's easier for them to separate from one another it's easier for them to enter the gas State and you'll end up with actually more molecules in the vapor phase and in the liquid phase so the weaker your attractive forces the faster your evaporation rate is going to me um liquids that evaporate very easily are called the volatile liquids liquids that do not evaporate easily are called non volatile liquids okay so here's a little summary of vaporization vaporization increases with increasing temperature with increasing surface area and with decreasing strength of intermolecular forces so the weaker our intermolecular forces the higher our rate of vaporization um let's look at this example here sum up which sample has the greatest rate of vaporization we have four 100 milliliter samples of water two of them are at 25 degrees Celsius two of them are at 55 degrees Celsius so right there you know you'd look at those um fifty-five degrees Celsius containers and you'd be like it's probably one of you guys and then we're gonna look at a hundred milliliter beaker with a five centimeter diameter and a 250 milliliter beaker with a seven centimeter diameter that's seven centimeter diameter is going to give us a larger surface area a larger surface area means that we'll have increased vaporization so choice D is our example that kind of meets both of our criteria when we're talking about the energy required for vaporization we do see that it takes energy for a molecule to separate from its liquid molecules and enter the gas state what that means is that vaporization is an endothermic process because it's going to require energy condensation on the other hand is an exothermic process okay so paper ization requires energy to overcome those attractions that energy requirement makes it endothermic the heat of vaporization therefore is always going to be an endothermic value and the heat of vaporization is going to be the amount of heat specifically that is required for us to vaporize one mole of any liquid okay that is called the heat of vaporization or the enthalpy of vaporization either term the heat of vaporization is always positive okay and the heat of vaporization is always going to be opposite the heat of condensation okay those are links and values they're linked by signs within any system we have what is called a dynamic equilibrium that becomes established all right so if we're talking about vapor and a liquid system if we have a closed container however much liquid vaporizes into gas form it's going to balance out with the amount of gas that re condenses into liquid form okay now there's never a point within your liquid that there is no vaporization and no condensation happening it happens all the time but the rate that it happens balances out so that there's no net change in the amount of liquid that is present in your container okay and that's the idea of dynamic equilibrium so we can visualize this so this is your initial container like you just poured in some water and you capped it and now you're gonna see what happens so initially the only thing that's going to happen is you're going to have liquid water that starts to vaporize okay now once you have liquid water that starts to vaporize as soon as you have water that starts to vaporize you're going to have vapor that starts to wreak on dense but you're still not at equilibrium yet so you're still going to have more vaporization happening then you have condensation happening because we still need to get to that perfect spot where we have equal rates eventually if you let enough time pass what happens is that your rate of vaporization so we can see these three up arrows here the rate of vaporization is going to equal the rate of condensation okay now at this point vaporization and condensation are happening at the same time and at exactly equal rates and this is the state of dynamic equilibrium that's when your your rates of opposing processes are exactly the same now that doesn't mean that you have the same amount of liquid in the same amount of vapor within your container it just means the rates that they are vaporizing and condensing are the same rates okay there's no net gain or net loss of material we can actually measure the pressure of a system when it's in dynamic equilibrium and that is called the vapor pressure so it's going to be the pressure that's exerted by a vapor when it's in dynamic equilibrium with a liquid so you can't really take a vapor pressure or something if it's not at equilibrium because it's not done changing it but once you're out of state of dynamic equilibrium there's gonna be no net change between the amount of liquid and solids so now you can take the pressure of the vapor over the liquid and get a can get a value for that as you might expect the weaker the attractive forces that are present among the molecules themselves the more molecules will be in vapor form as opposed to liquid form so the higher your vapor pressure will be so things that have very weak attractive forces have very high vapor pressures things that have very weak attractive forces are more volatile right because if you have a high vapor pressure that means that you have more molecules that are in gas form which means that you have a volatile liquid okay um if you change the volume of your container that's going to disrupt your equilibrium okay anytime that you disrupt your equilibrium it's going to take a minute but it will re-establish itself all right that's what the equilibrium will constantly try to re-establish itself with in a container so if the volume of your chamber is increased for example if you increase the volume think back to chapter 5 and your gases increasing volume will decrease the pressure of the vapor inside of your container when that happens well you're going to cause your condensation rate to slow because now you have to get back to the point you're gonna have to get back to that nice sweet spot between condensation and vaporization that you were at before you made the volume change so for a little while you're gonna have a rate of vaporization that's faster than your rate of condensation and that you know your vapor molecules will increase and then eventually your hit that equilibrium spot again and at that point your rate of condensation will increase so that it's the same exact rate as your rate of vaporization and equilibrium will be re-established at that point okay so this is just I mean this looks actually a lot like the water example before but in this case we can look at the changes that are happening here so we start with the system that's at equilibrium and now if we increase the volume we're going to increase our vaporization rate and decrease our condensation rate until we have a situation where they're at equilibrium again and then in that case our little arrows will be rebalanced umm if we decrease the volume we're going to increase our condensation rate and decrease our vaporization rate again until these things are at equilibrium again and then the arrow sizes would be the same our rate of vaporization and our rate of condensation would become the same once again okay so dynamic equilibrium is such an important concept in chemistry so just you know star this definition for yourself or something we spend a whole chapter talking about it in chemistry 104 um but it it basically says that when you have a system in equilibrium and it's disturbed the system is going to respond to minimize the disturbance and to return to a state of equilibrium okay so let's look at a conceptual connection what happens to the vapor pressure of a substance when its surface area is increased at a constant temperature okay if the surface area is increased well that's not gonna do it's gonna decrease the vapor pressure okay that's almost like increasing the volume if you increase the surface area you're increasing the space so you're decreasing the pressure vapor pressure is also dependent on temperature okay as your temperature increases your vapor pressure will increase because you will be moving more more of your molecules into gas form if you make very small changes in temperature these can result in very big changes in vapor pressure depending on the intermolecular forces that are present in your molecules okay so you can look at some of these vapor pressure curves um this is a typical curve it has the vapor pressure on the y axis plotted against the temperature on the x axis and you see the point here at 760 Torr or one atmosphere this is called the normal boiling point so I'm gonna define the normal boiling point for you in another slide or two but this is the temperature where you're basically where your liquids boiling at one atmosphere or at 760 Torr but you can see the the pressure changing right as the temperature increases as the temperature increases your pressure value goes up up up um when we talk about boiling I mean I know that I know what we think of when we think of boiling we think of like water in a pot you see the bubbles but that's not the technical definition of boiling all right the technical definition of boiling is the temperature of a liquid when it reaches the point where it's vapor pressure is the same as its external pressure okay so your vapor pressure and your external are going to be the same value um at that point you can have vapor bubbles form anywhere in your liquid okay so it's not just the surface of your liquid that can evaporate you are at the point where you can have these vapor bubbles forming at the surface of your liquid the middle of your liquid the bottom of your liquid alright so that's what's happening when you're actually seeing liquid boil this is a good example of it because you can see the the bubbles happening down here these vapor bubbles are at the same pressure as the external pressure when we say the term normal boiling point we're talking about the temperature at which the vapor pressure of a liquid is one atmosphere and the external pressure is one atmosphere excuse me um that's normal boiling point now we can have boiling points happen at different pressures right because if we have lower external pressures then we could have lower boiling points so for example if we wanted to look at different geographical locations that have different external pressures we could compare the boiling point of water at the top of Mount Everest which has a an external pressure of 0.32 atmospheres which is crazy low but the boiling point of water there goes down to 78 degrees Celsius okay and because it's it's it has it has to hit a lower threshold before it can vaporize if we have a place like Boston which is at sea level so the external pressure is one atmosphere well at one atmosphere you need the full 100 degrees before you're going to have the vapor pressure of the water be the same as the external pressure around it so we could see that that trend the lower the external pressure to lower the boiling point the higher the pressure the higher the boiling point we can also look at the heating curve of a liquid which is what happens when we're trying to heat a liquid to boil all right hmm regardless of the initial temperature that your liquid starts at if you're heating it you're going to see a linear increase in temperature okay so it's going to creep up up up up up up and your liquid will keep increasing in temperature until it hits 100 degrees Celsius at 100 degrees Celsius your water will not get any hotter okay your temperature will remain constant what happens at 100 degrees Celsius is that all the water that is present is going to convert into a gas all of that extra energy because you're gonna keep your flame on right it's not like you're gonna hit a hundred degrees Celsius and you're gonna turn off your stove flame and stop the water from boiling you keep that flame on you keep the water boiling well all that extra energy that you're inputting into the water it's being used to convert the water from liquid phase into gas phase and eventually all of your liquid will evaporate and all of your liquid will be in gas phase and say that you continued to heat the sample at that point now at this point once all of the water has been converted into water vapor the temperature of that vapor will now start to rise linearly again okay but for this period of time where you are where you have both liquid and the vapor present the temperature is going to maintain then I'm sure the water is going to maintain at 100 degrees Celsius until there is no water left so let's look at this question that relates pressure and boiling point use the following figure to estimate the boiling point of water at an external pressure of 200 Torr okay so this is our heating curve for water okay so if we want to get down to where about this about 200 it's just about 200 I would say it was close to the sixties ish right somewhere around there so one of our choices is 66 degrees Celsius so what that means is that if we had an external pressure of 200 Torr we can boil our water at 66 degrees Celsius at 66 degrees Celsius our vapor pressure of water would be the same as the external pressure of 200 Torr let's look at a math example here okay calculate the mass of water in grams that can be vaporized at its boiling point with 155 kilojoules of heat okay so with this in this example we're given the amount of heat that we have available and we want to know how much water we can vaporize with that so to answer this question we're going to need to use that heat of vaporization that we talked about earlier okay so the heat of vaporization is forty point seven kilojoules per mole well we know we started with 155 kilojoules so now we can relate 155 kilojoules to our heat of vaporization and we can figure out how many moles of water we could boil or we could vaporize if we had a hundred 555 kilojoules at our disposal once we know the mole value of water that we could vaporize then we can use the molar mass of water to figure out the mass of water we can vaporize so this is what this would look like mathematically we have 155 kilojoules of heat available to us this is our heat of vaporization so if kilojoules is oh is our starting unit and kilojoules will go on the bottom of our conversion and moles of water will go on the top so forty point seven kilojoules will vaporize one mole of water and every one mole of water there's eighteen point zero two grams which means that I can vaporize sixty eight point six grams of water okay we can also chart this relationship between pressure and temperature the best chart that we can make has a y-axis of the natural log of pressure versus inverse temperature in Kelvin okay and we can plot points along these two axes to give us a relationship between temperature and pressure but unfortunately that that graph it's it's not so easy to use if we don't have multiple data points a lot of times we only have a couple of measurements so we can reformat that graph relationship into a two points equation that requires only two pressure measurements or temper measurements temperature measurements for us to use okay it's not as precise as that graph that we just looked at but it's much more useful okay it also can let us predict the vapor pressure if we know the heat of vaporization and the normal boiling point among of a material because whenever we're told a normal boiling point we're being told a temperature and the corresponding pressure which has to be 760 Torr or one atmosphere because that's the definition of raw boiling point so this is the clausius-clapeyron three point equation and you guys need to know this equation and you have to know how to use this equation this is the the variables of the depth of the raishin to find it's the natural log of p1 over p2 equals the heat of vaporization divided by R and then in parentheses we have 1 over t2 minus 1 over T 1 ok p1 is going to be your vapor pressure at a particular temperature t1 in Kelvin so remember if p1 is over here and t1 has to be over here if p2 is over here then teach you has to be over here um I also know that this the version of the equation that I have written in this PowerPoint is slightly different than the one that's in your textbook and that's ok it's just it ends up being the same equation I've just modified the equation so that you don't have to worry about putting a negative sign over here ok because a lot of times when people are doing the equation that's in your textbook they forget about this negative sign here and excuse your answer in the end but if you use this version of the equation then you don't have to worry about the negative sign which makes a little bit easier I do remember that in this case your R constant is 8.314 joules over mole times Kelvin which means that your enthalpy of vaporization value that you're given has to be converted into joules before you can use it all right and let's look at an example in this of using this two-point equation so methanol has a normal boiling point of 64 point six degrees Celsius and a heat of vaporization of 35.2 kilojoules per mole what is the vapor pressure of methanol at 12 degrees Celsius okay so what are our first steps here we have two temperature values that are in Celsius if we're using the causes clapeyron equation those temperature values have to be in terms of Kelvin because our our constant is in terms of Kelvin um so we have to get rid of that we also need to let's write some of this down so we have to take these two temperatures and put them into Kelvin we also have this heat of vaporization value that's thirty five point two kilojoules per mole this value has to be converted in terms of joules okay Oh see this is the other version of the equation this is not the version I'm gonna use to solve it but I guess I forgot to change it in this one so let's switch over here and start to do our actual conversions so we go from sixty four point six to three thirty seven point eight Kelvin and we go from twelve degrees to two eighty five point two Calvin will convert our kilojoules into joules by just adding times ten to the negative three and the third thing we need to remember here is that at normal that the normal boiling point is at 760 torr okay so we were given this value even though they didn't expressly tell us what the pressure was because they said methanol has a normal boiling point of 64 point six degrees Celsius so they're saying at 64 point six degrees Celsius the boiling point is or the pressure is 760 Torr so let us plug some stuff in over here so we have Ln of P 1 over P 2 equals H rap over R and then our temperatures so we're gonna do Ln of 760 over P 2 equals thirty five point two times 10 to the 3 joules per mole over 8.314 joules or a mole times Kelvin and then in parentheses one over two eighty five point two minus one over three thirty seven point eight all right what we end up with is the ln of 760 over p2 equals two point three one okay so all of this stuff over here if you at all to our calculator correctly and ends up being 2.31 once we have our Ellen 760 over P 2 equals two point three one now we have to deal with separating out these terms some people like to take the e of this right at this point to get rid of the natural log and that's fine some people only do that I've taken the longer method just as just we could visualize it when we have the natural log of two terms if you want to separate those two terms you have to pull them apart with subtraction following your log rules so what this turns into is the ln of 760 torr minus the Ln of p2 equals two point three one at that point we can isolate unknown so we can isolate the Ln of P two I've added it to the right side and I've subtracted two point three one so I end up the Ln of 760 minus two point three one equals the Ln of P two and if you do the Ln of 760 minus two point three one that ends up being four point three two is equal to Ln of P two and at this point I take the e of both sides a of the natural log cancels out so I just left with P 2 and E raised to the four point three to is seventy five point four so seventy five point four Torr is our second pressure okay um it seems like a big you know big confusing formula but it's really not that bad just plug it right in there we're gonna start to look at some some other phase change points okay so we focused on the vaporization the condensation but other things are happening when we're looking at the phase changes between the different states we have what's called the critical point in a mixture and that's the temperature that's going to be required to produce a supercritical fluid okay and I'll define that for you on the next slide the pressure that the critical temperature happens at is called the critical pressure okay so obviously the critical point is the spot when you have your critical temperature and your critical pressure after that point it doesn't matter how much higher your temperature gets your gas can't be condensed to a liquid okay it doesn't matter the temperature and the pressure is gonna matter in that case what you end up with is what's called a supercritical fluid okay and these are gonna be fluids that have properties of both liquids and gases they are not a liquid or a gas okay the named supercritical fluid is a little bit misleading because it's not something that's in a liquid or the gas state it's something that is a weird hybrid of both right so as your liquids heated in your sealed container more vapor collects causing the vapor pressure to rise the density of the vapor increases the density of the liquid decreases and you'll actually see a meniscus between the liquid vapor the liquid and the vapor rather you're gonna see it you're gonna see it and then all of a sudden you're gonna stop seeing it okay so at this point when you're at your supercritical fluid there's no differentiation between your liquid state and your gas state anymore okay you have them basically it's almost like they formed some kind of homogenous solution with one another other state changes are while two other state changes are called sublimation and deposition and this has to do specifically with solids and gases okay and there's no middle change to liquid in either of these cases so sublimation is a change that's from a solid directly to a gas deposition is a reverse of that that is when you have a gas and it immediately forms into a solid okay this is one reason why molecular solids do have a vapor pressure because they can undergo sublimation and deposition okay if you have a if you have a molecular solid in a closed container and that molecular solid has a low enough low enough intermolecular forces and you can have this equilibrium form between a solid and the gas state so you'll have sublimation where the solid is forming into a gas I'm the Oh hug deposition where the gas is reforming into the solid but there is no middle point where those things would be in liquid form then we have fusion and fusion is a fancy word for melting okay so that happens when you have a solid that's heated eventually the particles that make up your solid are going to be vibrating to the point where they sort of shake free okay they're gonna have enough energy at some point where they are they will overcome the intermolecular attractions that are holding them in place and at that point they will have transitioned into the liquid state okay they'll kind of gain all of the properties that a liquid has like the ability to flow the ability to take the shape of the container that it's in so we call that melting or fusing okay but understand that these two things are the same process the opposite process of melting or fusing is freezing so that's when we're going from a liquid back to a solid and we're gonna look at this in a little more detail now just like we looked at the heating curve for the water now we're not gonna heating curve for a solid so if you have a solid at a particular temperature and you heat that solid okay again the temperature is going to increase linearly until you hit the diffusion point okay or the melting point of that liquid or that's all I'd rather um at that point your temperature is not going to increase okay the temperature of the solid is going to stay constant and it's going to continue to melt melt melt melt melt melt and at this point all of the solid has melted and turned into liquid now at this point once everything is turned into a liquid and you continue heating it now the temperature will begin to rise again and it will continue to rise until it hits the vaporization point in general should make sense melting is an exothermic process because melting is going to require heat or energy in order to have those solid molecules gain enough energy to separate from one another right that's gonna be like an energy cost having the molecules be able to separate from one another freezing on the other hand is an exothermic process okay that's not going to require any heat or any energy for freezing to happen the heat of fusion is called Delta H with a little F us it's an F of an endothermic process so that's going to be a positive process and the heat of fusion is going to be directly opposite to the heat of crystallization which is what we're calling that Delta H of freezing okay so freezing and crystallization those terms are going to kind of mean the same thing so we have a heat of fusion and a heat of crystallization and these two processes are going to be opposite your Delta H fusion value b-positive your heat of crystallization value will be negative okay and if you have a heat of sublimation that is going to be the sum of the heat of fusion plus the heat of vaporization because we're skipping that middle step with the liquid so these are just a comparison of heat of fusion versus heats of vaporization and you can see that those heats of vaporization are significantly higher than the heats of fusion now we're going to look we're gonna kind of put all this together we're gonna follow the heating curve of water you have to be able to look at a heating curve and do this calculation okay hold on I've seen Jirka water so the heating curve of water it follows the heating of water from the point where it's iced all the way through to the point that it's steamed okay and we're gonna be able to calculate the overall heat in input or energy input that was required to complete this process all right so five steps like I said and if we actually look at this chart the steps are are differentiated by whether they are increasing or whether they are plateaus okay so you see how we have three spots we were increasing and then we have two plateaus we can well let's go let's go through step by step I'll just tell you so segment one here that's an increasing step okay that is if we have 1 mole of ice okay this is our starting point we have one mole of ice and it's at negative 25 degrees Celsius um I'm gonna skip ahead but I'm gonna come back because I just want to show you the end point real quick now we are going to eventually take that one mole by and we are going to end up heating it to 125 degrees Celsius so that's that's the process that we're going through right now we're going from negative 25 degrees Celsius up to positive 125 degrees Celsius and we're going to track the energy change throughout that process so if you look back to this curve step one is down here this is the increase the first increasing line that we have so this is step one we're gonna track that he changed first then we're going to calculate the actual state change okay the solid to liquid state change that happens here our heat of fusion after we calculate our heat of fusion we're going to again track the change that happens while we increase the temperature after we hit our boiling point now we have to contend with the heat of vaporization after the heat of vaporization is dealt with then we can track the heat change that happened when we go from 100 degrees Celsius to 125 degrees Celsius okay so that's how it's broken down so segment one is we are going to take this one mole of ice which starts at negative 25 degrees Celsius and we are gonna heat it to zero degrees Celsius at zero degrees Celsius that is our fusion point so we know that we're not going to be able to increase the temperature of ice past one-half zero degrees Celsius the temperature of ice does not get warmer than zero degrees Celsius instead it fuses ok so we're going to do our Q equals MCAT equation from chapter six to figure out the amount of heat that was involved in making this transition so we know that the mass of one mole of ice is 18 grams that can be our mass the specific heat of the oh the specific heat of ice I got I just lost where I was in this um this 17 of ice is 2.09 joules per mole degrees Celsius we're not in the liquid phase yet and your specific heat values for ice water and steam are gonna be a little bit different so we take our mass multiply it by the specific heat of the ice and then we're gonna multiply it by the final temperature minus the initial temperature okay and that gives us our Q value and a Q value is a heat value that's 941 joules take that joule value and convert it right away into it into a kilojoule value okay so my Q my Q value for segment 1 it caught it uh it costs 0.9 for one kilojoules to turn that ice into liquid water okay or to raise it up to the temperature where now become liquid water because what happens in segment two is the actual state change this is the heat of fusion to figure out the amount of energy the particular fusion costs we have to take the moles of the material that we're looking at and multiply it by the heat of fusion the heat of fusion is in terms of kilojoules per mole that's why we are going to take the amount of moles that we are working with and we're going to multiply it by the heat of fusion to figure out the Q value for the process so the Q value for that state change between solid and liquid is six point zero two kilojoules so that's how much energy that cost okay and again it makes sense that these are all positive values because these are all endothermic processes so now our next segment is going to be us increasing the temperature from zero degrees Celsius now that we have them all this liquid to 100 degrees Celsius okay so we know we have we're back to our Q equals MCAT equation because we're tracking the temperature change anytime you're tracking a temperature change you need use two equals m cat because we have that part of our equation that accounts for our temperature change so Q equals M cat is 18 grams times the specific heat of water times the final temperature minus the initial temperature and we end up with 7.5 2 times 10 to the 3 joules which I forgot the unit in there apologies but that equals to 7.5 to kilojoules so that's how much energy it cost to raise our liquid water from 0 degrees Celsius to 100 degrees Celsius and now we're going to convert it into steam so this is our heat of vaporization we have 1 mole of water we're gonna multiply it by the heat of vaporization which is 40 point 7 kilojoules per mole and that costs us forty point seven kilojoules of heat okay now we have our last segment where we have our steam it's at 100 degrees Celsius and we are going to increase it to 125 degrees Celsius so now we're back to our Q equals M cat equation the mass is 18 grams the specific heat of steam is two point zero one joules per gram degree Celsius and our temperature change is 125 degree Celsius minus 100 degrees Celsius okay so this is gives us a Q value of 904 joules or point nine zero for kilojoules so if I wanted to know how much energy it took for that entire process now that I have each segment calculated well now I can add them all together so the overall Q or the overall heat or energy cost of this process is segment one plus segment two plus segment 3 plus 4 plus five if I add them all together I get an overall Q value or an overall energy cost of fifty six point one kilojoules to take one mole of ice at negative 25 degrees Celsius and raise it to 125 degrees Celsius in the steam forum sorry there should be a little negative sign in front of this 25 degrees Celsius here so that's the summary of those side steps that's a heat call that's the energy cost or the heat cost all right but that's the heating carbo water and you should be able to do that on your own okay let's look at this conceptual connection the heat capacity of ice is two point zero nine joules per gram degree Celsius the heat of fusion of ice is six point zero two kilojoules per mole when a small ice cube at negative 10 degrees Celsius is put into a cup of water at room temperature which of the following plays a greater role in cooling the liquid water warming the ice from negative 10 to zero from melting the ice okay um think about which of these things has greater which of these things has a greater has greater energy to it I think is the better way to look at it okay your heat of fusion which you know is your actual your melting point basically of the ice there are six point zero two kilojoules available in that so if you melt the ice you know you're going to be able to have more of an effect on the the temperature of the surroundings the next thing we want to look at our phase diagrams and these are going to be visual representations of the different states and the different changes that are happening at specific temperature and pressure conditions alright all phase diagrams are formatted in the same way okay there's going to be you know and I'll show you one in a minute so that you can get you could start to understand what they look like but you're going to have areas of solids liquids and gases on these phase diagrams and there'll be separated by lines and each of these lines is going to represent the state change that happens between those two phases um at the actual line both states would coexist okay the critical point on your phase diagram is going to be the furthest point along your vapor pressure curve there's also a triple point on your diagram and that is that the specific temperature and pressure where all three states of solid liquid and gas can coexist at the same time okay now for most substances the freezing point increases as pressure increases alright then you can use the phase diagrams to actually visualize this let's look at the phase diagram for water at first okay so we can look at actually at some of these points that we just talked about so our solid is always going to be furthest to the left our gas is always going to be furthest to the bottom right and our liquids always gonna be the top right okay so anytime you look at a phase diagram even if it's not labeled solid liquid and gas that's where you would find those different phases the triple point you see this is the area where all three lines meet so this is the point where you could have all three states existing at once the lines that separate the states represent the transitions between the two states so this line is called a fusion curve because it's the difference between the liquid and the solid the vaporization curve is between the liquid and the gas the sublimation curve is between the solid and the gas okay the critical point is the furthest point on the vaporization curve and pass this point this is where you would have your supercritical fluid okay here is another example of a phase diagram this is trying to tell you how you would navigate it okay so if you want to for example figure out what kind of I don't know if so if you were given a question I'm trying to figure out like how a question might be phrased if you're looking at a phase diagram so if you're asked what is the temperature or what state is this liquid in at one atmosphere in twenty-five degrees Celsius okay that's a valid question for a phase diagram then you would look at 25 degrees Celsius and you would trace it up until you hit the one atmosphere line and you could say all right well at one atmosphere in 25 degrees Celsius this particular material is in its liquid state if I asked you what state is the material in it negative 25 degrees Celsius and 1 atmosphere well that's over here so it would be in a solid state then okay what else can we see from these things we can also look at the C the tilt of this line here that tells us the relative density of the liquid to the solid material okay so for example as I increase my pressure okay in something like iodine the more I increase my pressure the more my the more solid I'm gonna get right so my fusion curve here it tilts to the right when your fusion curve tilts to the right like this that means that your solid phase is more dense than your liquid phase okay so in carbon dioxide 2 you see we have a very slight tilt to the right so if I was in Crete it's harder to see on this one because it's more subtle but if I were as over here and I was increasing the pressure eventually I would intersect my solid phase that means that my solids more dense than my liquid now if I look back at the water diagram will look what happens here what I'm increasing my pressure I am NOT hitting my solid phase okay it's easiest to see at a point where you could potentially intersect both of them so if I look at 0 degree Celsius and I increase my pressure increase my pressure the first thing I hit is my solid phase and if I continue to increase my pressure I'm back into my liquid phase so my fusion curve tilts to the left for water when my fusion curves hilts to the left like this my solid is less dense than my liquid so you can actually look at a phase diagram and tell by the tilt of the fusion curve whether you have a solid material that is more or less dense than your liquid so in the case of water with this tilt to the left you have solid that is less dense than the liquid something like iodine and carbon dioxide that fusion curve tilts to the right it is more dense than the liquid okay so let's look at a conceptual connection here a substance has a triple point negative twenty four point five degrees Celsius and 225 millimeters of mercury what is most likely to happen to a solid sample of the substance as it is warmed from negative 35 degrees Celsius to 100 degrees Celsius at a pressure of 220 millimeters of mercury okay so this is trying this question is asking you to imagine what the phase diagram of this substance would look like now if the triple point of the substance is at 225 degrees or 225 millimeters of mercury rather this is indicating the pressure level that you would have to reach before you were able to transfer into your liquid state okay so the fact that we're keeping our pressure below that it kind of takes one of our states out of contention okay a triple point at negative twenty four point five degrees Celsius that means that once we move past once we take this thirty five negative thirty five degrees Celsius solid and we move it to a temperature that's higher than negative twenty four point five degrees Celsius that solid can no longer be in the solid phase it's going to transition into a different phase the pressure is not high enough for it to transition into the liquid phase so it must sublime directly into a gas phase okay the last two slides that we have to look at or more a summary of the intermolecular force conversation that we've already had why is water such like a weird substance it's a liquid at room temperature where most things that are the same molar masses water and our molecular substances or gases at room temperature okay but not water water is a liquid again hydrogen bonding right we have a lot of hydrogen bonding here the strong intermolecular forces keeps the water in the liquid phase water is an excellent solvent it's able to dissolve things that are ionic and polar because of its large dipole moment it has a very high specific heat for something that is a molecular substance and this is again due to its strong intermolecular forces practically speaking for us that's why it's able to have a good temperature effect on our coastal climates right that high specific heat allows us you know it's able to like absorb some of that temperature change that occurs in the environment also water expands when it freezes at a pressure of one atmosphere and this expansion is going to be one of the things that makes it less dense than liquid water okay so it's one of the few substances that has a liquid state that is more dense than its solid state but that's about it for this chapter let me know if you have any questions just shoot me an email