we are alive I think if you guys can hear me see me you know if all of that you know technological stuff is working to drop in the chat and there may be a one through five as to how you're feeling about AP chem how everything's going so far you know just wanna make sure you guys can hear me and see where we're at in terms of you know AP chem photos hang on a minute wait for a couple more people to kind of shuffle in I know we kind of got a smaller crowd tonight kind of hanging out for a few minutes tonight we're just gonna be going over unit three so it goes through it goes through the things like you never like there are forces and solids and gases and really a lot of the stuff involving gases which is really important to your AP exam and then a couple other kind of miscellaneous topics like solutions and other little things so I think maybe we'll hang out for like another minute or two let people shuffling shuffle in they want to and if not we'll just jump right in all right guess we'll just jump right into it and just share my screen here there we go all right so first things first you know be sure to follow findable and all this social media things you know Twitter Instagram YouTube all that jazz alright so here's a little outline of what we're really gonna be going over the first thing we're doing is intermolecular forces which is kind of the whole branch off point for this entire for this entire unit everything that really am me talking to you you can kind of connect back to intermolecular forces so it's a really important topic not just for this unit but for the AP exam as a whole it's really something they love to ask you about because they're relatively simple but if you don't know them you really don't know them so tell you that's really important to understand involving the AP exam then we're gonna go into solids liquids and gases and more of the in-depth things regarding them so you know how they interact how they act the different types of solids stuff like that then the ideal gas law and kinetic molecular theory which are all things involving gases and we're also gonna be talking about deviation from the ideal gas law they're gonna be talking about solutions and solubility and mixtures and then finally we're gonna end it off with a couple little things involving light and how things like things that absorption work which we'll get into later with the beer-lambert law so this is jump right into it so this is the first thing we're gonna be talking about is our intermolecular forces right and these are really important to your AP exam right they they love to ask these on just any type of free response question because they could easily just shove in you know what air molecular forces are found in this and boom there's a whole part right there and so it's really something that is important to know for AP exam so really intermolecular forces are just simply put attractive forces between molecules so if you for molecule here molecule here they feel some sort of attraction that's an intermolecular force right and there are four main types of intramolecular forces right so they're ion ion forces ion dipole forces dipole dipole dipole forces and London dispersion forces and we'll go through each for these in detail and then the ones with dipole dipole and London dispersions those are called Vander Waal forces you really need to know that but it's kind of just something that if they do refer to it that's what they're referring to so the first is the first is the first type of intermolecular force that we're gonna be looking at tonight our ion ion forces they're pretty much the simplest and you've seen them before right they're basically just ionic bonds their attraction between oppositely charged ions and they're the strongest in a regular force right you've got a positive ion here a negative ion here opposite charges they you know opposites attract and so they're held together by the electromagnetic force opposing charges so the reason that they're the strongest is because rather than being held by something like polarity where they're kind of partial charges or or other tensile forces with ion ion force they're straight positive negative charges right and they're held together proportionately to Coulomb's law which we learned about in the last unit right that that is based off of the distance of the ions and the charges of them themselves that leads to a stronger interaction so ion ions to ions will have the strongest intermolecular force and that kind of explains why when you see things like salt right like table salt it's so hard to melt because if you want to break those ion ion bonds you're gonna have to put in a lot a lot of energy right so that's why assault has a boiling point like a thousand degrees Celsius because you need a ton of energy to break those intermolecular forces and turn it into liquid salt all right to the next one we're gonna be talking about our on dipole forces and again these you've probably seen in your class before right they're just attraction between an ion and a polar molecule so we see them but most common example and really the the only example I can think of would be something dissolving in water right so you see we have these polar waters and we'll go into polarity a little bit later surrounding the positive and negative ions right so we have an interaction between the the negatively charged or positively charged ion and the positive or negative sides of the polar molecule in this water so we see here we have water right we have a negative sign and so it's attracted to that positive ion and we see the water's surround it right so with the negative side facing the sodium similarly we look at a negative chloride ion we see that the waters are flipped right because that now the positive end is attracted to the negative ion and so that's why you've probably seen a picture of something like that before that's where that actually comes from from the ion dipole forces between some sort of ionic compound and water or another polar solvent typically all right so the next one and arguably the most important one Aurora called dipole-dipole forces so dipole-dipole forces are really just simply put attractions between polar molecules all right so you see this polar molecule here we know that because it's polar it will have a positive end and a negative end right and so because of this when we have multiple molecules you know floating around together the negative end of that molecule in this case the chlorine right will be attracted to the positive and end of the other in this case hydrogen we see the same thing happened with iodine monochloride I think that's called right we see the negative chloride be attracted to the positive iodine in this case so that's really how dipole-dipole forces work and they they're really common and they've always found it literally any polar substance right so ammonia water hydrochloric acid hydrofluoric acid you know chloromethane any sort of polar molecule will have dipole-dipole forces right and they are the relatively strong they're not in it really depends on the molecules the how strong they are but there is a form of dipole-dipole forces called hydrogen bonding and hydrogen bonding they're the strongest form of dipole-dipole forces right so really they're the same the Hyderabadi tends to really trip up students right and the reason for that is that they think that it's different from dipole-dipole forces what a reality these are the exact same as dipole-dipole forces however they're stronger because they're between hydrogen and a highly electronegative atom which the AP exam considers nitrogen oxygen and fluorine to be the ones that you will see exhibit hydrogen bonding so we see here with water right we have the negative oxygen attract to the positive hydrogen so this is just a dipole-dipole force however what makes it special is the fact that we have a an intermolecular force a dipole-dipole force between the hydrogen and the oxygen making it a hydrogen bond which is stronger than a normal dipole-dipole force however they are the exact same right this is the same as this bond this that bond this form from before except just stronger due to the fact that oxygen and hydrogen will form a hydrogen bond right same thing happens with ammonia right if you have to ammonia ions or to ammonia molecules we see this this magnet attracted to the hydrogen because of the polarity of ammonia and because it's an N bonded attracted to an H we see again a strong hydrogen bonds form and I know we see a couple people kind of happenin now how many equals about this I know I'm moving kind of quickly feel free to drop any questions you have whatsoever even if it's not really super related to this I'm here for you guys so if you have any questions whatsoever you know feel free to drop them in the chat and I'll be kind of popping in back and forth from daya from time to time after the last type of intermolecular force so we're gonna be looking at are what are called London dispersion forces and there are three types but I think of the evening exam I'm fairly certain you can just refer them as London dispersion forces and so when we're looking at letting dispersion forces but they have to do with nonpolar molecules right so you might be thinking to yourself okay got a non-polar molecule so no forces are felt between the molecules however that's false because of the fact that if we're dealing with a non-polar molecule we're still going to see that due to just randomness right random movement of electrons around this molecule they're going to what are called induced dipoles right so they're going to be very instantaneous and very weak dipoles found in nonpolar molecules so we see here for example that it just for an instant is is pulled over to a negative side and pulled over to a positive side and so that induces a dipole so in a perfectly nonpolar situation we have something like methane for example we're going to see incredibly weak induced dipole forces because methane is nonpolar if we see something like methane mixed with a polar molecule we're gonna see dipole induced dipole because we see the induced dipoles of that non-polar molecule interacting with the dipole of a of a polar molecule and the same thing can be said for an ion induced dipole situation or if you have a charged man such as we have here we're going to see that attracted to the negative side of this induced dipole of a non-polar molecule there's kind of a very tricky situation kind of get through your head and it takes a little bit of practice of looking at but essentially induced dipoles are found in every single molecule whether you're polar nonpolar and ion induced dipoles will always be found except for in ionic compounds I kind of said if you are whether you're polar or nonpolar you're going to feel induced dipoles because of the fact that because the fact that you know the random movement of electrons is still going to create some of it all right so when we're talking about induced dipoles you know you might be wondering well how do we measure the strength of these induced dipoles right so we use a measure called polarizability and I know there's kind of a big word but all it means is just how easy is it to induce a dipole basically how you know how often are we going to see these random movements of electrons causing a dipole situation so typically just the rule is that the bigger an atom is the more polarizable it is and therefore the stronger the induced dipoles or it's for example if we're looking at iodine and fluorine right you look at your periodic table and you can see well iodine is way way way way bigger than fluorine right and so by that we can make the conclusion that intermolecular forces with an iodine are going to be stronger than that in flooring even though they're both going to be induced dipoles and in fact if you look at the periodic table and that some charts you'll see that iodine at standard temperature is solid whereas fluorines at gas so that's why we see differences in intermolecular forces even within nonpolar molecules it's because of this concept of polarizability right because of the fact that certain atoms will have better abilities to actually induce a dipole than others all right so here's kind of the definitive right people this is really the thing that people really want it's the definitive list of just you know what are the strongest to the weakest IMF's strengths right so we see that we have ion ion right all the way at the top then ion dipole then we have hydrogen bonding I kind of gotta fix this then we have there we go then we see hydrogen bombing right then finally we see dipole-dipole forces and London dispersion forces so that's really the list in order and some anything to memorize this but it just kind of makes sense that ions are gonna be stronger than dipoles and that dipoles are gonna be stronger than induced dipoles right it kind of makes sense logically but it's really not that hard also to kind of just memorize and it becomes second nature after doing a couple of practice problems so for example here is an example question right there aren't going to be many questions that just straight up ask you which has the strongest intermolecular forces they'll ask you to use the concept of intermolecular forces to answer a question so for example this one is which of the following most likely has the highest melting point right so when you think about what what the melting point really is because you're looking at what is the strongest really to actually which is the strongest to break those intermolecular forces right that's what you want to do in this situation and so with this question you want to look at well basically which of these has the strongest intermolecular forces so the first thing we can do is we can immediately knock out ch4 because of the fact that it's nonpolar right methane is a non-polar molecule and you can know that just by drawing out the Lewis dot structure of it you'll see immediately that it's nonpolar and we can knock it out with hydrofluoric acid it is polar you'll note that however when we're looking at B and D you'll note that both of those are polar and exhibit hydrogen bonding right so really I kind of wrote this question poorly but it would come down to whatever the strongest intermolecular forces so for example if choice B was changed to something like co2 right now the answer would immediately be h2o right because of the fact that carbon dioxide is nonpolar methane is nonpolar hydrofluoric acid does not exhibit hydrogen bonding and D does exhibit hydrogen bonding which is the strongest intermolecular that is shown out of these four molecules and so that's really how you kind of want to look at intermolecular forces there really aren't gonna be many situations where they do ask you you know which of the following has the highest intermolecular forces they would ask something about melting point or boiling point or or something like vapor pressure which we'll go over later and using those qualitative aspects in order to extrapolate data using intermolecular forces right so do you have any questions on that I know we kind of blasted right through that but it is one of the most important concepts in this unit so I just want to make sure you guys really understand it what h2o still have a higher melting point than nh3 in terms of the AP test they wouldn't ask you that because they both have the same intermolecular forces I can check I know them we know the melting point of water right let's see melting point of h2o is zero degrees Celsius right whereas nh3 is yeah water would have a higher melting point than nh3 due to the fact that oxygens more electronegative than nitrogen right so we're gonna see stronger intermolecular forces in h2o we're gonna see stronger hydrogen bonding in h2o compared to nh3 does that make sense yeah sounds good right really when you're doing these problems and really when you're when you're taking the AP test in general the periodic table is gonna be your best friend I'll drop a quick website that I have to bookmarked actually when I took AP chem it's an online periodic table that I'm sure most of you guys have heard of but it's honestly your best friend when you're doing these types of problems because you can immediately look at it and say okay this is clear from just the chart right alright so let's move on to the next topic right so I kind of combined two the key concepts here into one topic and so that's on solids liquids and gases that's right so we all know from elementary school right the three main states of matter right solids liquids and gases and we know the main difference it's right if I put a cup of water in front of you you would know oh that's a liquid anything can we a little bit of a step forward and kind of try to describe the behavior of solids liquids and gases so what were we talking about with solids are primarily the main types of solids that we're gonna be looking at so the four main types that I'll be going over tonight and that will be on your AP exam our ion excelence covalent Network solids molecular solids and metallic solids right and all four of these might sound really intimidating really difficult but they all have very distinct characteristics that makes sense when you once you understand how they work right so like before the simplest had to know are just ionic solids right so again with ionic solids we're going to see ions attracted to each other look at a most like a cross-section of an ionic compound will see that they form a lattice right a crystal lattice of positive and negative ions in order to reduce reduce repulsion and my is traction right we see that that's sodium ion right here for example this sodium ion is attracted to four different chlorines and vice versa chlorine is attracted to a bunch of sodium's and so you see that we because of the strength of these intermolecular forces we're gonna see that an ionic solid is going to be brittle right now whenever I think of an ionic solid I always think of just a bar of pure salt right a bar pure salt is party it's hard you'd have to snap it in half and so it's brittle right and that's because of this tightly packed crystal lattice that's formed between the ions right it also has a high melting point right as we discussed before those ionic forces are gonna be the strongest type of intermolecular forces and so it is incredibly difficult to break apart those imf's they don't actually conduct electricity in the solid phase however when they're dissolved right remember electrolytes and non electrolytes from previous chapters when they when an ionic compound is dissolved in water it's broken up into its ions right when you dissolve salt and water you get sodium ions and chloride ions you're going to see that we get we get a solution that conducts electricity due to the presence of these right so this is pretty simple early is the simplest type of solid because it's typically the one most students are most familiar with it's moving on to covalent Network solids right so there are two types there are connected networks and there are layered networks so looking at a connected network first the way I like to think of a connected covalent Network solid is just a giant molecule right so when we're looking at something like diamond for example diamond is pure carbon right however what makes what makes diamond so hard so rigid and gives it the properties that has is the fact that each of these little lines here is actually a covalent bond between carbons right so really what a diamond is it's just one giant mega molecule of carbons all covalently bonded together so covalent bonds are the strongest type of interaction there are right I know we were talking about ion ion forces however covalent bonds are literal bonds right they are genuine tried and true bonds whereas with ionic compounds we're going to see intermolecular forces holding together the ions where the covalent bond interesting literal sharing of electrons bridged ions bridge atoms right sure connection and so we see that this is the strongest type of talk trucks stronger than a solid right so with a diamond for example you can't break in time and it's the literal hardest substance on earth similarly when we're looking at something like graphite right graphics it's the same as diamonds pure carbon what gives it different properties it's actually the way it is a solid so when we're looking at graphite we see that these are covalently bonded right these little lines are all covalent bonds however however when we're looking at those covalent bonds we're not seeing that there's no covalent bond here we see sheets of covalent bonded layers of carbon connected by London dispersion forces and remember London dispersion forces are the weakest type of intermolecular force so that's why when you put your pencil of graph on the paper it just shreds right off the pencil because of the fact that those London dispersion forces are so weak so you can break the graphite right so the major thing you should take away from covalent Network solids is just basically that they are rigid and hard right you know graphite still rigid you can only snap it in half and they do not conduct electricity there there's no way to get those electrons moving right here they're all rigidly bonded together here they're all rigidly bonded together and so we are not going to see any sort of conductivity with something like diamond or graphite or coal all right there's a lot of carbon things with covalent Network solids but we're not gonna see conductivity whatsoever either the third type this picture might look really similar to what we were talking about before however with molecular solids these aren't actually covalent bonds these lines between these water molecules are actually intermolecular forces in this case hydrogen bonds and so when we're looking at something like solid water or just ice we see that the water molecules aren't held together by any sort of bond any sort of ionic force but they're held together by weaker intermolecular forces so these weaker intermolecular forces give it they get softer and they give it lower melting points right so I always like to kind of have it an example for these regular solids I think of ice right and you might think well ice has a pretty low melting point right because of the fact that these hydrogen bonds happen it will melt in your ice will melt in your hand right all you gonna do there 232 degrees whereas something like salt if you want to melt salt you gotta get into thousands of degrees because you have to break those intermolecular forces right and finally we'll have we have Mitama type metallic metallic solids right so metallic solids are they're just metals right but the thing with metals is that you have ad localization of electrons right so you have the positive nuclei of this of this atom the same iron or something right so metal and instead of having you know electrons here electrons here electrons here we see what's called an electron see where electrons are free to move around them and this is what gives metals its major properties right there are the greatest conductors of heat and electricity right that's because these electrons are given free rein right so no matter in the solid phase liquid phase you're going to have conductivity and heat and electricity right that's why we build our bill that's why we make like frying pans out of metal because they conduct that heat so well think about if you made a you know think about if you made it pan out of something like you know foam right ignore the fact that it would melt right it also wouldn't conduct the heat all that well similarly metals are malleable right they're able to be bent and they're ductile which means there's we bent - bent into wire right so things like copper are really great for making wires which is really convenient because we take a metal which is able to conduct electricity then Andrew wire and boom we've got the ability to use electricity so it's really something important about metallic solids the final thing with this is just the difference between crystalline and forest solids I'm gonna go through this really quick it's like one little bullet point in the curriculum basically crystalline substances have this nice grid structure so this is something like you know something like diamond right whereas this is an M forest solid it doesn't crystal lattice type look to it write something like glass or sodium I'm studying or sulfur dioxide right or yeah I think it's Oprah dioxide I'm not sure actually silicon dioxide that's what sand is so that's the difference between a crystalline and an M for a solid so finally keep same finally we're gonna look at the energy of solids liquids and gases and this is something that even if you haven't formally learned it you kind of know right a solid is really rigid right so if you want to make it a liquid you weaken those intermolecular forces and you put in energy into the system and you get this this liquid right then you want to break the intermolecular forces and create a gas by adding more energy right no get more into this we need for actually looking at things like thermal chemistry and thermodynamics however it's really something that is kind of something to look at here where you have to put in energy and break those intermolecular bonds or intermolecular forces all right so does anyone have any questions on what we were talking about just with types of solids or anything to do I'll just kind of hang out exciting cuz I know I'm moving kind of quickly it's just cuz there's a ton of information in this unit alright any questions you guys can still feel free to ask any questions about what I'm going over from the past I'm just gonna keep moving forward all right so the the site the next topic that we're gonna be talking about is the ideal gas law right this is like the Big Kahuna of gases right this is basically going to be applied to if not every I'm still in Neenah - and I'm chillin hey that's great yeah chillin here to learn random stuff that's awesome man you know it's also way you get to you two three you're gonna kill him and you have all this information you're heading even in the back of your head are you killing it right the next thing we're gonna be talking about is the ideal gas law right so the ideal gas law looks at how we can relate things like the pressure the volume and the temperature and the moles of gas together in one really frankly a pretty beautiful equation right you might not think it's very beautiful but when you start using it trust me it gets really convenient so in my opinion and in probably in most people's opinions the ideal gas law is the most important equation within this course regarding gases and we're gonna be going over variations in the ideal gas law so we will you know don't we will be going over so right just to go over what you know PV equals NRT actually means right so it says the pressure times the volume is equal to the moles of gas and the gas constant gas constant yes R times the temperature so the thing that actually people tend to mess up with the ideal gas law isn't the math it's the unit's so for example they might give you this and what are called Torr or millimeters of mercury milliliters mercury and then they might give you this gas constant in atmospheres right you have to make sure and I cannot stress this enough you have to make sure all of units are order right we don't really have time tonight to go over specific problems involving the ideal gas law unless you guys want me to in which case drop a question a little bit later I'll be happy to go through some but yeah really the most important is just making sure your units are in order right especially are are there a hundred million different values for R there's 0.08206 liter atmosphere from prim okay remember I know there's 8.314 Joule per mol K I think they're a bunch let me see if I can actually find them for you it comes through there see there's like five really it's more convenient to actually convert your units and use 0.08 206 because that's just the first one they give you but yeah they give you it 4 joules atmosphere and tour so you do have different values to deal with so if you have to or want to you could also swap out are just the important thing is to make sure that all of your units line up right if you have atmospheres for pressure you gotta have atmospheres in your gas constant if you have leaders in your volume you got to make sure you have leaders in your in your gas constant right Kelvin for temperature Oh ease use Kelvin for temperature you should never be using Celsius when you're dealing with these always convert to to kill them and to do that you just add 273 to your Celsius so that's really the ideal gas law in a nutshell it's just really an equation that once you get some practice with it once you learn how to use it it becomes second nature right they can only ask you so many questions involving PV equals NRT but it's super convenient because then if you get a pressure a volume and like a gas constant and temperature you can immediately find out moles all right and go from there so it's a very convenient equation so the nice thing we're gonna be talking about is dalton's law of partial pressures right so again this is a topic that really tends to screw up allows students but really important simple term is when you have a mixture of gases the full pressure is the sum of the partial pressures so right so when we're looking at if you have a mixture of let's say three different gases right let's call it a B and C the total pressure is just going to be the pressure of a plus the pressure of B plus the pressure of C right and you do that so on and so forth for however many gases you have so really that's the main idea and there are other equations that I'm showing you here that involve finding the actual partial pressures but that's when you're given the total pressure really this X a is a pretty important topic it's called the mole fraction right it's just basically the moles of a so how many moles you have this divided by the total moles it's just a matter of finding a percentage of how many moles make up the total it's fairly simple in terms of the actual formula right and these are all they're given to you right so on the aps and they give you PV equals NRT they give you I rip these actually directly from the the curriculum so they are going to be given to you but it's really simple and easy to memorize especially because you will be applying these pretty heavily alright so here a couple graphs involving gases that you may have to know there are two main gas laws outside of the ideal gas law that you might want to know and that are pretty convenient right so the first one is Charles's law so what Charles's law talks about is basically the relationship between temperature and volume right so according to this graph right as we increase temperature we increase the volume right think of this like a hot air balloon rate we increase the temperature of the gas the gas expands blowing up the hot air balloon right and that's why hot air balloons work because of Charles's law similarly Boyle's law looks at the relationship between volume and pressure right so as you compress a gas as you add more pressure you are going to decrease the volume right at a low pressure we're gonna high volume at a high pressure we have a low volume right does that make sense I know these graphs might be a little bit confusing at first so I'm just gonna make sure I got like a thumbs up from a few people just making sure that y'all are getting this hey Bert awesome right yeah this stuff is pretty cool I mean it helps you understand this is really the chemistry that you really understand the world right you look at a hot air balloon you're like oh look at that Charles is long and then your entire friend group is like dude shut up about chemistry already right so here's another alternative form of Boyle's law so here's the graph on the left that we just saw before right on a high pressure yeah i pressure right we're going to see yes a high pressure right we're gonna see a low volume and at low pressure we have a high volume right similarly if we graph pressure versus one over the volume we get a linear relationship right and this is because of the fact that this is an inverse relationship and so we invert the volume we get a linear relationship we see that at right at a high pressure we had a really low volume and sort of high pressure we're gonna have a really high inverse of volume it's really kind of confusing it's I still kind of to wrap my head around it but just know that if you see this on a test this is just Boyle's law this is just Boyle's law okay a linear relationship between pressure and the inverse of volume so what we can use these the these gas laws right Charles's law is Charles's law Boyle's law and what's called gay lussac's law right Charles law Boyles long gay lussac's law which relates temperature to pressure and volume we've called the combined gas law right I know you're saying oh my god not another equation this one you actually should memorize I made the dumb mistake of not memorizing it and not remembering to use it and barely managing to get a question right on the AP because of the things that I understood the ideal gas law this is something that's incredibly stupidly useful it's used basically to just change calculate changes when other things are held constant so this actually something where I do want to go through a problem with you guys and kind of go step-by-step through how to actually do a problem so here I'm going to pull up this little online whiteboard thing that I've been using for the past couple streams and I'm going to there we go where we gone I have to just get the actual problem into the whiteboard there we go right so to say that assuming a constant temperature right so constant temperature if one liter if a one liter container of gas at a pressure of 0.32 atmospheres is constrained to 0.25 liters what is the new pressure all right so we look at the the we're going to look at the combined gas law which states that p1 v1 over t1 is equal to p2 v2 over t2 right however we know that temperature is constant so what we can do is just get rid of the temperatures right they won't impact what what we're looking at here so what this boils to that boils Havilah is p1 v1 is equal to p2 v2 right and so from here what we want to do is just plug in our values right so our original pressure is 0.32 and our original volume is 1 liter is equal to p2 we don't know even though the new volume is 0.25 so that ok so then all we do is isolate solve for p2 by dividing by 0.25 and then from there you would just plug into a calculator right so that's really the way that you're going to be using the combined gas law and the one caveat that I really don't like with the combined gas law is that they don't give it to you actually they will not give you the combined gas law on the formula sheet at least they didn't give it to my class and so I don't want to say they will give it to yours for something that you really want to memorize but it kind of makes sense by understanding Boyle's law and Charles's law and gay lussac's law all right let me pull up the slides again right so that kind of brings us to the end of our equations involving gases right so now we're gonna be talking early why does this happen right why does pressure exist why do certain laws exist and that's what's called kinetic molecular theory comes into play right so there are a few assumptions that the kinetic molecular theory makes so the kinetic molecular theory is a theory that helps describe why gases act the way they do so for assumptions and I kind of bolded the important parts are that gas moves in random directions that collisions have no energy losses that they move in straight lines until colliding with the container and that the average kinetic energy of a sample of gas is proportional to temperature so let's break those four down for just a quick second I kind of threw a bunch of information at you at once so let's break these down breaks of gas moves in random directions at relatively large distances so basically what this says is that when you look at a sample of gas it's just going right we don't have any order really to the actual molecules we just see them flying every which way in random directions and we see and we assume that they're relatively large distances from each other so we're not gonna see them bumping into each other we're not going to see interactions between the molecules similarly collisions are perfectly elastic right so we see if we have a container and a molecule bumps off of it we assume that there is no energy loss right it's not going to slow down the molecule we also make the assumption that gases move in straight lines that kind of makes sense right we don't see a gas under the kinetic molecular theory go like this where we don't speak like that and then right we just see boom boom straight lines and then the one that I consider the most important is that the average kinetic energy the sample of gas is proportional to temperature essentially what this means is that as you increase the temperature your molecules will get more kinetic energy which means they will move faster right again kind of make sense you put energy into the system and they're gonna start moving more and faster right we can we'll be looking at something I think it's on the next slide of were called Boltzmann diagrams involving how we look at distributions of gases right and their kinetic energies right so we look at gases and kinetic energy we know the equation kinetic energy is equal to 1/2 MV squared where V is the velocity we don't you don't need to apply this on the exam but it's in the curriculum in terms of understanding it this is why when you increase the temperature you had an increase in average kinetic energy because they're going to increase their velocity right and therefore more energy or that's why when you increase temperature you get an increase in speed because when you put energy into the system you get an output of velocity right so this is really the the conclusion from the relationship between kinetic energy and speed right so what this diagram shows it's not actually showing the speeds of any specific molecules rather it is showing the distribution of molecules that are moving at a certain speed so we see that at this super low speed not many molecules right but if we look at this speed right in the middle we see the most number of molecules are at that right so this is what we call the most probable speed right and there are other speeds like the average speed right here but really we what you really want to understand is that that this curve does not show the actual speed it shows the distribution of speed so if you had something like this and I'm kinda gonna draw this really bad like this even though the curve is lower it's moving faster because we have a higher distribution going a faster speed right in the AP exam parlance I'm really sorry that really sucks actually we're trying to get our own platform up and running but right now we're kind of stuck with chromecast so it's kind of you know a love-hate relationship right so this is what maxwell-boltzmann diagrams show right so let's try to look at a practice problem right so we see which of the following Boltzmann distributions are at a higher temperature right so remember that a higher temperature implies higher kinetic energy so really what this question is asking is which of the following gases are are moving faster so is it the orange one or is it the green one and I'll tell you if you can't really see it that well this higher one is the orange one and the lower one is the green one so parth says green let's see if anyone else has any different answers or if everyone agrees all right all right give it another minute or two seems that has a higher average perfect right that's perfect Green is the correct answer because although we see that there isn't peak just quite like the orange the Green has a higher average has a higher average speed right we see that more of the molecules are moving at it's like 500 something compared to 200 for orange so that's a typical problem that students make they see this big peak and they think whoo high speed when you're looking at the axes you see that this is the number of molecules or in this case actually it's the probability but yeah it's the number of molecules right pov yes it's the probability of velocity so basically yeah the number of molecules the higher you are here the greater probability it is that you're moving at that speed so yeah the more molecules there are that's really how Boltzmann diagrams really work they're essentially statistical distributions of probability as opposed to but it's helpful to think of them as just a number of molecules right and you'll never be asked to graph one like this however there was a question last year I'll try to find it where they asked you to draw a Boltzmann diagram after a temperature increase so that's a question we might want to look at if I can find it real fast I'll give us like 10 seconds to try to find it if we don't you know too bad yeah now I can't find it but I think it's in 2017 so if you happen to find it give it a try it's a really good question for this unit in terms of looking at the the Boltzmann distribution oh there it is I found it right we don't have time to do it though so yeah I'll just drop it in here it is number one Part B do number two like bii a temperature increase do we make the peak higher and compress the graph with a temperature increase you'll see the curve flatten out right so actually let's let's do this problem right just for the sake of making sure you understand what that means in terms of a temperature increase so the context of this problem was that the temperature was decreased actually I think we'll have to read the question but let's take a look all right no you're good parts don't worry about it that's actually a really good question that you asked because they very well could ask about two different curves and the differences between them so it was a really good question so don't don't worry about it all right so here we see the graph below shows a distribution for the collision energies of reactant molecules at a 120 degrees Celsius right so it says draw the second curve on the graph that shows the distribution for collision energies of reactant molecules at 30 so you might be thinking you know Dylan what the heck this isn't gases however you can treat it the same way right this the energy of collisions can be really said as kinetic energy right or velocity worth and fraction of collisions well that's just the same thing as the number of molecules right a higher fraction so a higher number so if we want to draw at a lower temperature right excuse me it's like trying to find it and I lost it there we go I got it right so if we want to draw it at a lower temperature we want it to be at eight we want a higher percentage to be at a lower energy so we're going to see something that looks like this right we see a higher percentage at a lower temperature compared or a lower energy compared to the other right and then we're gonna see it did below this function cuz a lower percentage are at a higher a higher speed so that makes sense to how you would draw different lowering or hiring the temperature if we made the temperature higher well what we would do is we would actually draw something like this it would kind of flatten out and you would just see a higher proportion at a hydrogen so that makes sense parth I know that was your question so lower temperature means less movement so how can that make more collision okay so this is kind of getting into kinetics which we'll get into later essentially we're not looking at collisions we're looking at the energy of collisions right and the fraction collisions the number of collisions hypothetically would remain the same and so if we draw this lower temperature right like that I draw that again there we go that looks good right we see that there any lower temperature so at lower energy so we see that a higher fraction of of collisions happen at this lower energy compared to when we have a higher temperature that make sense we're not seeing actually more collisions we're seeing more collisions take place at a lower energy we're seeing the same amount of collisions actually take place according to the bolt from diagram awesome alright and this dream probably I'm guessing is gonna go above an hour because there's they're like 13 points in this actual unit so it's gonna take a hot minute to get through but if you guys have any questions feel free to stop me asked me to go through a problem right I'm here for you as I said a couple of times right I'm here to help you guys understand the material not just blast through it as quickly as humanly possible so I know parth you said you were on unit 2 if you have a question about unit - be my guest man drop a question great so the next question I know we had one person who was asking about this is gonna be deviation from the ideal gas law right the ideal gas law is actually complete utter fiction there's no such thing as an ideal gas there gases are very very close to an ideal gas however an ideal gas actually doesn't exist right every gas deviates from the ideal gas law so what we see is that as pressure increases right we're going to see deviations from the ideal gas law so when we look at these graphs we see that we have PV over n RT right you might think you know what is what does that mean but remember for an ideal gas PV is equal to NRT so PV over NRT is just 1 and so that's why we have a graph of 1 right here and any pressure according to an ideal gas this PV over an RT will be 1 however looking at gas as different temperatures right we see that they they deviate from the ideal gas law right so when do really gases differ from the ideal gas law PV equals NRT so there are two there are two scenarios right high pressure so a high pressure gas molecule start to take up space right so the actual assumptions that there are a negligible distance for each other it goes away right the kinetic molecular theory so they start to take up space compared to the actual pressure right because when pressure increases remember I think it's Charles's law volume decreases right or no it's I think it's boils actually isn't intermolecular forces will start to take effect right where to low temperature we start slowing these molecules down and suddenly they start feeling that attraction to each other right because they have less energy to go flying around so this is an equation that actually corrects for these two factors and it's called the Vander Waal equation so really this might freak you out first of all I'd like to say that you will never have to actually plug into this equation for ap chemistry there will never be a situation on the AP exam where you'll be plugging into this you just have to understand it right and thinking about this forever so this term has to do with the high pressure term right so when volumes really small right or when volume is actually negligibly large this is going to approach zero similarly with this it'll also approach zero and so for ideal gas PV is still equal to NRT but for a non-ideal gas PV is not equal to NRT it's going to be slightly more or less cover the ideal gas law is still really handy if we are at a regular pressure and a regular temperature due to the fact that these terms will actually kind of cancel out and they won't have as much of an effect on the end answer so PV equals NRT is essentially a simplification of the van der Waal equation we're two here a couple little things that have to do with it so kind of looking at this right we see that those four molecules kind of get tightened in right I have to kind of restart the slide to get the actual animation to keep going right assumes that the space is available for all the molecules to move in excuse me but only this amount of room is available because the rest is taken up by the molecules themselves right and that has to do with the actual deviations from the ideal gas law similarly in a real gas we see attractive forces that occurred maybe see you know if we have water vapor we're still going to see those hydrogen bonds happen and so that's where the ideal gas law fales so I know there was one person who was kind of iffy on it so does it make better sense now or do you still have a couple questions on it all right looks like it for okay kind of keep going on right so our next topic is mixtures and solutions right so the first thing we learned is calculations involving molarity so molarity is a measure of concentration basically how much how much solute right how much stuff do we have dissolved in how much solution so molarity is the most used one on the AP exam and in really chemistry in general molarity basically is moles per liter if you have for example four moles of sodium chloride dissolved in three liters of three liters of solution you have a 4/3 molar solution so say the same thing here right if you have a three moles of hydrochloric acid dissolved in three liters of water you have a concentration of 1 molar right 3 moles per 3 liter which is one mole per liter which is one molar and the notation for this isn't really relevant now but it will be heavily relevant later is when you see something in brackets that means concentrations specifically it means molarity right so another application of molarity will be dilution right decreasing concentration by adding water right it's always it's always gonna be water and you want to use the equation m1 v1 is equal to m2 v2 right it makes sense because you need the moles of solute remain the same you're not adding and you saw you and so you need the moles per liter times liter to equal the moles per liter times new leader right and so that's how dilution works right and it's pretty simple to understand when you need to use it right just a change in molarity or a change in volume the next thing is representations of solutions right so when we're looking at a solution remember they're gonna be two things there's the solvent and the solute and the solvent dissolves the solute so we see what actually happens is the solvent and the solute expand and form a solution or they directly form a solution but at the end you end up with this form of a solution where you have the cell you integrated in the solvent and that's really the whole point of of a solution right you have mixed thoroughly throughout you have a homogeneous mixture right it's kind of a weird thing that if you're a visual person it makes sense in terms of just come is trivial to visual people and really it is kind of symbol in terms of right makes sense you have the salt that you dissolve in the solvent right so now we're gonna be talking about how do you separate a solution right if I have something dissolved in something else or having a solution to liquids for example how do I actually separate them so the first new one is called distillation so distillation takes advantage of boiling points and differences in vapor pressures in order to separate a solution so this may look like a really complicated thing that's going on here but really what's happening is we have I think a hot plate here right and what happens is we set it at a specific temperature to boil the lowest boiling of the two things in the solution right so for example if you have I don't know two two liquids and we have one that boils at 70 degrees one that boils at 1700 right we set it at 70 degrees and that first bit will start to dissolve or will start to not dissolve will start to boil and so it vibrates out of the solution and it goes into this long tube labeled five so what happens is when it goes into that tube we use typically water to cool it down back into the liquid phase and collected right here right so it's a pretty simple process at its base level there's some really complicated distillation that I've seen but really that's really the main point you just you boil it you have different boiling points so you boil one out is me mainly the main point the second one is called paper chromatography so chromatography separates mixtures by taking advantage of polarity and non polarity right so for example when you've an ink spot right it's a mixture actually of a bunch of different inks and so if you put a dot on paper and you can do this at home if you put a dot of ink on a little piece of paper and you dip it in water the water will travel up the paper and pull the ink up the paper based on the polarity right yeah most most Kevin classes end up doing some sort of paper chromatography I actually did an eighth grade I didn't do it in AP chem but it's something that yeah a lot of people have been exposed to and it's really totally simple okay so moving on we're gonna be talking about solubility right but it is kind of weird seeing the water go up right it kind of travels up it's pretty interesting them right so solubility is just what it sounds like so this looks like a lot of text thrown at you at once the solubility rules that the AP test wants you to know specifically group 1a elements nitrate and ammonium are soluble those are the ones you really need to memorize however don't dwell on memorization it becomes pretty clear what's all you'll and what's not through usage right you can memorize your solubility rules if you really wanted to but typically you learn what is and what isn't by just doing it you learned the silver chloride is highly insoluble you learn that silver nitrate is highly soluble that sodium you know that ammonium chloride is soluble something like that and the major thing to take out of this is that like dissolves like a polar solvent will dissolve a polar solute etc etc right so really this is just a list of rules solubility has to do with it has to do with basically how easily something can be dissolved and you'll find later that technically everything is dissolved a little bit but it's based on these rules basically show you what does dissolve a lot right what will actually dissolve in water to an extent that you can see it as moving on we see we're gonna be looking at the electromagnetic spectrum and the photoelectric so kind of a big jump from what we were just talking about so the electromagnetic spectrum I think most of us have probably heard of it right it has to do with wavelengths of light right remember according to quantum theory light is a wave right it is it has a particle wave duality and based off of where you where a particle of light exorbitant light exists on the electromagnetic spectrum is based off of this wavelength right and it goes gamma x-ray ultraviolet visible infrared microwave radio I don't know if you actually memorize it I don't think you do but basically what it has to do with is high frequency high frequency or a right or short wavelength short wavelength suddenly to longer wavelength right or lower frequency right and that's basically the trend across the electromagnetic spectrum and we'll be looking at basically what it means to have a low frequency a high frequency stuff like that so there are a few equations for light that you will need to know so H is defined as Planck's constant right and at six point six to six times 10 to the negative 34 joule-seconds so really what Planck's constant shows is that light is quantized right when you put a little part of light energy will be released in quanta little packets that are proportional to the frequency so basically this brings up the formula thought the energy of a photon is equal to Planck's constant times mu mu what is that H nu right that's the Greek letter nu that is the energy joule seconds is kind of a weird a really weird unit kind of ignore the units in terms of their actual unit units just know that thing but it this way right yeah it's a little weird right so when we multiplied joule-seconds times so the unit for frequency is per seconds right basically how many how many you know wavelengths do you see per second and so you see the per seconds cancel the seconds and you're left with joules right so looking at this image we see that for example when an electron moves down a couple of energy levels it releases it releases a photon of light in the form of energy proportional to H times the frequency right so that's really where it comes from and here's another image of it right that energy is a minute in quantum right little packets of H nu right so E is equal to H nu ya the Balmer series is I believe from the second to the second energy level I believe so second about the speed of light can be defined by wavelength that's that little grief letter lambda and nu right we see that C is equal to land to do where C is three times 10 to the eighth so the unit's really do make this clear right meters per second is equal to meters for the wavelength times per second for the for the frequency so meters per second will give us the speed of light which is has improved by Einstein to be a constant right so basically just looking at those two equations so the second part is looking at what's called the photoelectric effect so essentially what the photoelectric effect is is that when you blast a piece of metal with light you sometimes see an electron pop off right you see boo boo right and you and the real problem was it was solved by Albert Einstein such the one thing he won a Nobel Prize for was for solving the photoelectric effect right he figured out yes evey is different rules B is electron volts it is a measure of elastic an electron volt is a unit of energy which is it's module regs one electron volt is it's really iffy it has a really scientific key thing but just know they're different from jewels so really what einstein and other scientists and whether other chemists and physicists had actually observed is that when you blast a piece of potassium or other metal with light we see an electron sometimes pop off sometimes not so you kind of ask okay well why right it's because when you add enough energy to it you're going you have to add enough energy to knock off an electron right to promote the electron all the way off of the metal and so when you use something like red light it does not have enough energy actually to knock off an electron whereas when you use something like green or purple light you do you do knock off an electron you can you can go into the physics of this and the quantum mechanics all day but really the basic idea is just blast it with light and electron is knocked off so finally the final thing we're gonna be talking about tonight is beers law of absorption so there's gonna be a picture after this but basically beers law is is can be summed up by the equation a is equal to ABC the absorption yes there is a stream on this in two days right so the person can be going to this in much more detail so I'm gonna go through this kind of quickly a is equal to ABC basically the absorption is equal to the molar absorptivity which is kind of a constant times the path length on flame that is times the concentration of the solute so if used actually find concentrations so it's used in a type of experiment called it's called spectral spectrophotometry right so for example we do this lab last year in my chem class you will find the concentration of red dye in gatorade right so you blast it with green light and you want to see how much green light will be absorbed right so here's a little picture that might make it clear right when you blast light through a solution right some of its going to be absorbed and then the rest is heavy Tran minute and so assuming a constant B or in a constant path length which is right here this is the path length right we see that a is equal to M or a B times C so we see a direct relationship between absorption and concentration fo really go more into that because I know there is a stream on this as as Perth mentioned so yeah so that is unit 3 right along last we made at the end I'm going to chew up for another minute or two just you know do you guys have any more questions anything else to mention questions comments concerns right I know I kind of blasted all the way through that material because there's a lot of it so if you guys have anything that you're really confused on feel free to ask right now you know it's not they know this end up blowing off I have for another thirty to thirty seconds or so see if you guys have any more questions good stream it's good thank you yeah hopefully you know feel free to share this with as many people as you can yeah I'll be streaming I I mostly do the reviews so exist end of stuff the big hour-long like mega reviews but I'm happy to if you guys want to do some problem videos or some other streams on random other topics but yeah I think on that note I'm going to log off for the night I will see you guys next time stream which is will be sometime in the future and I will talk to you guys later