 I just wanted to remind you guys that we have a lab practical this week. So instead of going to lab, well, instead of doing a lab, you're going to go to lab and do a lab practical, okay? So hopefully you guys have, so I posted all the quizzes and all that stuff, all we've done in lab on the lab blackboard site. So you should have everything that you need in order to do that. Remember exam twos this Friday. I think we'll cover up to all of next time. I'll talk more about that at the end of class today. I've posted two practice exams, hopefully you guys have been trying to do those already. I'm going to post those solutions probably sometimes today. Okay, so by tonight you'll probably be able to see. Remember exam reviews tonight, six to eight p.m. and Tuesday after this class, or after my other class, sorry, 9.30 to 11. But one tonight is going to be in this room, HSC 201. The one tomorrow, I don't know, I'll let you guys know on blackboard as to where this is. Okay, so are there any questions about any of that stuff? So lab practical, I'm just going to give you a note card and you're just going to write answers on the note card. If it says do a calculation then you'll do your calculation. Remember the same thing, box your answers and all of that stuff. Okay, so other than that it looks like nobody's got any questions so let's keep going with chapter six. So we're talking about this kinetic molecular theory and dynamic equilibrium. So this is the process at the molecular level where the same amount of product is converting to reactant and reactant is converting to product. So remember, those pictures really do do a good description of what's going on. But let's look at the second one of those two pictures. So the first picture this shows, you know, at time zero let's put the liquid into an evacuated container. So there's no gas up here. Since there's no gas up there, the liquid particles are going to move to occupy that space. And what you'll notice is that over time the liquid decreases in volume. It's because some of those liquid particles, liquid molecules became gaseous molecules, of course. So gas takes up the entire volume of the container. So it doesn't matter how many particles there are, right? It's only this amount. But it's taking up this whole volume now. But once that happens, once you've got what you would refer to as saturation in the gaseous phase, what happens is that the exact same amount of molecules that are going from the gas to the liquid are going from the liquid to the gas. So you have an equal proportion of these things going back and forth. We call this a dynamic equilibrium. And what you'll find is that a dynamic equilibrium sets up every time you've got two different phases of the same molecule in a closed container. So if you're, I don't know, boiling water in a closed container, you'll have a dynamic equilibrium of that until it explodes, you know, until something pops up or whatever. So any sort of system will behave this way until the phase. So let's talk about the gaseous phase. So we've talked about the gaseous phase. So the gaseous phase is characterized by low density. Of course, density is mass per unit volume. So remember, gas is very light. So the mass is very light. And the volume is very big. So they're very, very low in density. So an indefinite shape, that just means it takes on the shape of the container. So if you were to ask what's the shape of the air in this room, you would say the shape of the room. So it's an irregular shape or an indefinite shape that depends on the shape of the container. It has a very large compressibility. You can squish gasses down very well easily. That's why you can take a big volume of gas and put them into those little pressurized containers. It's because there's a lot of distance between the actual particles of a gas. So when we're looking at a solid, remember, a solid, if we're looking at the individual particles of solid, they're stuck together like this. So if I were to want to squish this, I couldn't squish it very much because there's not very much space in between those particles. A liquid is a little bit less kind of like that. They're still touching each other, but there's a little bit more space, whatever, so you can squish it a little bit. And then a gas, if you think about a gas, there's a lot of space between those particles, so it can compress them very easily. Take that space. So thermal expansion, that's if you have a very cool gas. You have a very cool gas. What you'll see is, if you cool the gas down, you may see some sort of mist on the bottom of the container, and then you heat it up and it'll expand out of the container. So it'll replace this space in between those particles. And you can see that happening here. All of those things. Low density, obviously low density. You have a container in both of these. This spigot right here was open, and then it allowed the gas to diffuse to the other container here. Large compressibility, of course, you can compress this amount of gas to half that size, obviously. That's fairly large when you compare it to solids, right? How much can you compress? You could never compress a solid to half that size. And then thermal expansion, I guess this doesn't really show that very well. Okay, so let's talk about ideal gases. Okay, ideal gases are model gases. And the way gases are supposed to behave, or ideally behave, of course there's no such thing as an ideal gas. It's like any sort of ideal whatever. It's not something that exists, but something that you would think about attaining to. Okay, so it's a theoretical gas made up of what we call randomly moving point particles. In a model of the way that gases behave at the microscopic level. Okay, so even though this doesn't really exist, it's good to think about, and you can start making up equations about ideal gases, and they really correlate to what gases behave like naturally. Okay, so natural gases behave like. So when we're thinking about the ideal gas concept, we can measure the following variables of gas, the temperature of the gas, the volume of the gas, the pressure of the gas, and the mass of the gas. So we want to remember that no truly ideal gas exists, but noble gases and simple non-polar diatomic gases like fluorine, fluorine, bromine, well fluorine is fluorine I guess, oxygen, nitrogen, and all the noble gases, they behave very very similar to what we would describe as an ideal gas. So very nearly ideal behavior at ordinary temperatures and pressures. Okay, of course if we put a lot of pressure on these particular substances, so we've got thousands and thousands of atmospheres of pressure on them, or we cool them down very very low temperatures, or heat them up to very very high temperatures, they don't always behave ideally, but it's still a useful model. Okay, so these are the things that we're going to be emphasizing, temperature, volume, pressure, and mass. So what happens is we can systematically change one of these properties and what will happen is the other three will change. Okay, so we'll see the effect on all the other three. So we want to talk about, so remember kinetic molecular theory, let's think about ideal gas connect. Okay, so this isn't, again this is ideal stuff. Okay, so gases are made up of small particles that are constant random motion. The distance of separation is very large compared to the size of the individual particles, so what that means is gases are mostly empty space. Again, why you can compress them very easily. All gas particles behave independently. Okay, so this is the first attribute that doesn't refer to known gases, only ideal gases. Okay, so they say no attractive or repulsive forces exist between them. Of course all substances have attractive and repulsive forces, but we're talking about ideal gas. Okay, so gas particles collide with each other and the walls of the container without losing energy. Of course when you collide with things you're going to lose a little bit of energy. Again this is talking about ideal gases. And the average kinetic energy of the particles increases or decreases in proportion to the absolute temperature. So what does that mean as the temperature goes up, the particle speed goes up. So you get more excited more flowing when you increase your temperature. Okay, so again gases are easily compressible this is why when you go to the fair or whatever and you see the clown like spinning the balloons, like twisting them up into animals, like here, it's because there's mostly empty space. You can twist them like that. So room for the particles to be pushed together. Gases will expand to any available volume. You know that by looking at a hot air balloon. Gases have low density. Also you can talk about hot air balloon. The hot air balloon floats away. Low density. So things with lower density are on top of things with higher density. So the gases in a hot air balloon are actually lower in density than the air around them. That's why they float away. So the gases have low mass per unit volume. They readily diffuse through each other. It's because there's so much space. So all gases you can mix all gases together readily. So if I wanted to mix the green and the white gas here because there's so much space between them that they're not kind of repelling each other. There's no like polar non-polar stuff going on. Gases exert pressure on their containers. You probably know this before and after you open like a two liter bottle of soda. You can feel the carbon dioxide that's in the soda bottle prior to opening it very pressure up. So it makes the bottle very stiff. Of course when you open it for the first time it becomes very, you know, less stiff. Very much less stiff. It's because all that gas was released. The pressure was released because the gas was released. So this pressure results from collisions of gas particles with the container of the walls and they're constantly colliding back and forth and giving this pressure. So there are no ideal gases. Again gases which behave most ideally are non-polar molecules at low pressure and high temperature. Okay so we've been talking about temperature and we remember temperature hopefully remember how to convert from temperature scales whenever you're talking about ideal gases and using these gas laws that we're going to be talking about all of your temperatures have to be in the kelvin scale. Okay so it's the absolute temperature scale and that's the temperature scale that molecules and atoms and other particles behave according to. Of course farenheit and Celsius are kind of made up, temperature scales, humans made up. The kelvin one is again an absolute temperature scale. So the temperature of a gas sample is the measurement of its average kinetic energy. Remember the hotter the sample is, the more energetic the particles are. Okay so kelvin temperature scale like I said is going to be used in all of our calculations in this chapter. Remember absolute zero is the temperature of zero K, zero kelvin. This is the temperature which all molecular motion ceases. Okay so that means that the average kinetic energy is zero. Okay so at this temperature kinetic energy becomes zero because all motion stops. And remember on the Celsius scale absolute zero is negative 273 degrees Celsius. So you may be having to convert back and forth from farenheit, Celsius, and kelvin. Okay but remember always you're going to be using kelvin in these temperature calculations. This here is an equation to relate temperature and kinetic energy. You guys don't have to worry about it. Okay we'll go back to pressure, talk a little bit more about pressure. So pressure again is the force pushing on a unit area of the surface which the force acts. So whatever that means right. So it's just these little balls hitting against the walls of this container is giving this thing pressure. It's the force exerted by the collision of particles within the walls of the container. It's often expressed in units related to the measurement of atmospheric pressure. Okay so what you feel on top of this right now by the air is called atmospheric pressure. And especially since we live at about sea level right we're getting about one ATM on us right now. So ATM is the unit of pressure that you're going to see most often and it's the unit of pressure that you're going to have to convert your pressure units to in order to use it in the ideal gas laws. Okay so now we've got two measurements that we need. So for temperature we want to convert that to Kelvin. Pressure that converts to ATM. You can see standard atmospheric pressure. One ATM is the pressure needed. This is another way of looking at it. It's the pressure needed to support 760 millimeter of column of mercury in the barometer. So what happens here is if you've got a pool of mercury and you invert a glass tube that's closed on one end, what will happen if you stick that inverted tube into your glass dish here? The pressure of the atmosphere will push down on that liquid in the dish and make the column of mercury go up that inverted tube. So what happens here at sea level what you'll find is the height of this tube here, the height of the mercury in that tube is 760 millimeters. And that's where if you've ever heard of this pressure readings, millimeters of mercury, this is where that comes from. So this is, you know, the barometric pressure a lot of times is given in tor or millimeters of mercury. So one tor is the pressure needed to support a one millimeter of column of mercury in a barometer. So the pressure of 760 tors equals one ATM. So pressure of one tor, the column would be up to there. So the pressure of one tor is like one 760th of atmospheric pressure. So again, as Earth gravitational field pulls gases in the air towards the center of the Earth, this is what's happening, this is what gives us the pressure. It results in a force that's on all objects of 14.7 pounds per square inch. So that's also one ATM. So 14.7 pounds per square inch equals one ATM. 760 millimeters of mercury equals one ATM. 760 tor equals one ATM. So tor and millimeters of mercury are the same thing. Tor is just some guy's name, tor, he TORR, he's long since dead now and he did a lot of measurements of mercury. So this instrument that we've created here, this mercury dish with an inverted 2-minute, it's known as a barometer. So it's the instrument used for measuring relative atmospheric pressure and like we said, the mercury goes to the height of about 760 millimeters above the dish at sea level, like what we're at about right now. Okay, so what you can do is, let's see, the diameter of the barometer doesn't matter. As long as you have enough mercury within the dish, you can have a fat barometer or a bare skin barometer, it'll go up to the same height. Of course, if you didn't have enough mercury in that dish it wouldn't, but say you did have an infinite amount of mercury. So the ratio of the weight to area always stays constant, so the mercury will always rise to 760 millimeters, and the pressure is inversely proportional with the height. So the higher you go, the less pressure you've got. Okay, so what you see here is the height of 0 meters, so sea level, right, you get the pressure of 760 millimeters of mercury. But if you go up to the top of a very tall mountain, 8,848 feet high, you see that the pressure drops dramatically 760 millimeters of mercury. Okay, so here's some more units of pressure that you may have heard or seen. PSI is also a pound per square inch and you can see the ones that we want to be concentrating on are atmospheres, and millimeters of mercury, but just to convert that. So what you can find is that if you've got the density of two liquids and the height of one liquid, you can figure out what the height of the other liquid would be. Right, so if we knew that at 0.5 grams per mil, the atmospheric pressure at 1 atm, what's the height of the mercury column we'll get to? Does anybody remember? 760. So if we go height of the mercury column at 760 millimeters, what would the height of a water column be? So we could figure that out by doing this equation. Who would the height of a water column be? So you've got a, if liquid one is mercury, and liquid two is water, so we'd be looking for the height of, well, we'll make this liquid one. The height of water, right, would be, and if you figure that out, it would be a pretty tall, used mercury in ornace, okay? And it's still a liquid, so maybe you want to try to figure that out on your own. Okay, so if you recall diffusion, diffusion is the movement of one gas or another. This is due to the particle nature of gases and the extreme distances between the particles. Gas can freely go into and diffuse in between each other. So this is why when you smell perfume, you can smell it, okay? It's because the perfume has vaporized and, you know, floated around the room colliding with the other air particles in the room, the nitrogen and the oxygen in the carbon dioxide, until it landed on your nose and you can actually smell it when it comes into contact with you. So remember particles of substances must actually come into contact with your nose when you detect them. So if you smell something good, it's because it's in your nose. If you smell something bad, it's also in your nose. Yeah, so if you think about bad smells. So if you recall this experiment that we've already done Ammonia diffused 12.8 centimeters in 3,760 seconds and Ammonia diffused 69.85 centimeters in the same time right? So if you notice, this is much smaller distance than this, okay? So about 5 times as small. And if you look here you'll see that HDL and NH3 Ammonia and hydrochloric acid, right? Ammonia is a much smaller molecule, okay? So it's going to move faster. This is actually what happens is that if you've got smaller molecules, they move faster. Lighter molecules move slower. You can think about this like by trying to compare throwing like a golf ball up in the air relative to a bowling ball up in the air, you know? Which one would be easier to throw around? It's because of the relative weight of the two particles. So infusion as opposed to diffusion, okay? So infusion is the process by which individual molecules flow through a hole without infusions between the molecules. So it's like these ducks in a row, right? So they're all flowing through the same particle, are all flowing through the same core, but they're not contacting each other. Thanks so you can see it is maybe better here. Okay? So a grand flaw of infusion, for gases that equal temperature and pressure flowing through one or another or through a different solvent, it deserves that the rate of one gas over the rate of the other is equal to the square root of the molecular or the molar mass of one over the other. So again, this is similar to this, right? If I have the rate of gas one, the rate of gas two, you should be able or if I have the rate of gas one and know the molar mass of both of them, you should be able to find the rate of the second gas, okay? So these rates will be given to you in like meters per second or something like that, okay? So we could think of an example, right? Let's go back to these. So say, well we know the molar mass of HDL is said that the rate of one of them, I don't know, is 10,000 meters per second, so that would be the, we'll say the rate of, what's the rate of HDL? You should be able to figure that out. This is a fusion and remember it's different than diffusion, okay? So when you go back and look at these numbers relative to how far they went, centimeters per second, it won't add up, okay? So if you try to do this, it won't obey this law here, okay? Because this is diffusion, not diffusion. I just wanted to make sure that we understood why this experiment was giving us the numbers that it was, okay? And that they were different than this law here. Okay, so that's all the lead up really to the gas laws. So I really want you to be able to do this diffusion stuff and this density to height stuff, okay? I really do. Let's talk about the gas laws now, okay? So gas laws are mathematical equations, so we're going to be doing a lot of these different equations with a lot of different variables starting now, okay? So the gas laws are mathematical equations that describe the behavior of gases as they're mixed, subjected to pressure or temperature changes or allowed to diffuse. So gas laws involve a number of relationships between number of moles, volume, temperature, and pressure. So you've got to remember how to calculate moles from grams and all of that, okay? This is why we concentrated so much on that earlier in this experiment because we're going to be giving this really hard. Yeah, the pressure and temperature are very... So the gas laws, here's three of the gas laws. One of them is called oil's law, Charles law, Gay-Lusak's law. There's another one called Avogadro's law as well. And then there's the combined gas law and the ideal gas law as well. But these are three out of the six that you'll know by the end of this chapter. So again, they're mathematical equations relating pressure, temperature, and volume of gases, and they're named after the scientists who first discovered them. So notice Boyle's law is the measure or the relation between pressure and volume. Charles law is between pressure and temperature. Gay-Lusak, or volume in temperature. Gay-Lusak is between pressure and temperature. And Avogadro's is between number of moles and volume. So there's Boyle, I'm a dead guy, obviously. So what he came up with was that the volume occupied by a gas is inversely proportional to the external pressure. So if you have these at constant temperature and number of moles of gas, so if you don't add or take gas away, the volume will be inversely proportional to the temperature. And that makes sense, right? Because, or the pressure. Because pressure kind of keeps volume in. Okay, so if you're squishing something down, right, you're going to make, you've got a lot of pressure on something that's going to be very small, right? That makes sense. So pressure and volume are inversely proportional. So you can, this again is like these types of problems, okay? So let's try this one together. The gas at 10 liters, occupies 10 liters at 1 ATM. If we double the pressure to 2 ATM, what's the volume? Let's figure this out together. Initial and F means final. Where it says VI, PI, VF, VF. That means initial and final. So hopefully everybody can see that the initial volume's going to be 10 liters. What would be the initial pressure? We've got the final volume. So we've got PF, right? So if we're looking for VF, all we've got to do is divide both sides by PF. Yes, PF cancels there. And we get VF equals VI times PI over PF. Cancels with ATM. So we get 10 times 1 is 10 divided by 2 is 5, right? So 5.0. You're going to be doing, like I said, those other two are similar to these, except the one has the square root sign. So you've got to watch out for that. Charles Law, Avogadro's Law and Gay-Lucic's Law, they're all very similar to this. So here's a couple that I'd like you to try on your own. Notice some of these give you more information than you need, right? You don't need the temperature for Boyle's Law. So this one's giving you the temperature. Remember the right number of significant figures. So let's talk about Charles Law now. So the Charles, John Charles is French guy. He was the balloonist. He wasn't even a chemist, right? But he figured out that the volume of a gas is proportional to the temperature. So if he heated something up, it would give big, just like a hot air balloon. So let's look at Charles Law. So his law is this VI over TI equals VF over TS. Now remember T's always have to be in Kelvin. So if your temperature is in Celsius, you're going to have to convert it to Kelvin. This one says a gas occupies 10 liters, so that's going to be what? Of these variables, 10 liters. VI at 273 Kelvin, what's one side going to be? TI and then doubling the temperature, 546 Kelvin that's going to be TI? Well, we don't know. This is upper TI. So if we're looking for VF, just multiply both sides by TS. A lot of these. So if you're wondering how do you solve these things, go watch these office hour videos. They must have 30. So we just take TI, 46 Kelvin, 273 times 10. There's some Charles Law. Now notice these are in Celsius, right? We're going to have to convert those to Kelvin. Gay-Lusak's law is the pressure of a gas is proportional to the absolute temperature of a gas. There's Gay-Lusak. So PI over TI equals PF over TS. So very similar to Charles Law. And you guys can try it on your own here. Again, if you guys want to do these during the review tonight or tomorrow, it's no problem. But I'd like you to try this one on your own. So you're going to have to remember these formulas. These five formulas that we've gone over. In fact, we're going to learn a couple more before the test. The combined gas law is derived from a combination of Boyle's Law, Charles Law, and Gay-Lusak's Law. So what you'll find is if you squish all of those laws together, you can make this law, which is the combined gas law, PI times TI over TI equals PF times PF over TI. Subscript I, you might also see subscripts of 1 and 2. That's the more antiquated way of talking about initial and final. So they'll say 1 is initial and 2 is final. So if you see those, don't be freaked out about it. It just means that initial and final. Here you can see combined gas law problem. It's set up for you. You can see, again, we need to convert our degrees Celsius to Kelvin. Notice our pressures are in ATM. If our pressure were in millimeters of mercury, we would have to convert that to ATM. If we were in tor, we'd have to convert that to ATM. I'll give you the conversion factor for pounds per square inch to ATM if I wanted you to do that. 4 millimeters of mercury you should know. 760 to 1. 4 millimeters of mercury are the same thing. 764 equals 180. Here's another combined gas law problem. Another combined gas law problem. And then the last thing we'll talk about is Avogadro's law. Avogadro's law is equal volumes of gas contain equal numbers of molecules. So the volume occupied by gas is proportional to the number of moles. So notice it's very similar to Charles' law. Instead it's VI over NI equals VI over N, or very similar to Gay-Lusek's law. They're all very similar to each other. It's just these four laws kind of correlating two of these four variables, pressure, volume, temperature, and number of moles. Remember, number of moles you can get if you're given the molecular weight and the number of grams. So you're going to be given that a lot. You're going to be given the mass number of the stuff, the mass. You're going to have to convert it to moles and then do this gas law problem. And for Avogadro's law a very interesting thing that you'll figure out is that you'll have a mole or volume because all moles like a mole of hydrogen contains the same number of particles as a mole of bromine does. That makes sense. They both contain 6.022 times 10 to the 23rd particles per mole. Any mole of all the platinum contains. So what you find is that the volume is directly proportional to the number of particles in the sample. What you'll find is that if that's the case, then any sample that contains a mole of particles of gas particles will be the same volume. And at standard temperature and pressure of course you can't be varying the temperature and pressure. So if you say the standard temperature of pressure is 0 degrees celsius and 1 atm, the volume occupied by one mole of gas is always going to be 22.4 liters. So that's a useful conversion factor to remember is that the molar volume is the one mole equals 22.4 liters. I'll write that down. Just to emphasize it's very useful. One mole of gas at standard temperature pressure equals 22.4 liters. The temperature and pressure are switched. There's something else and it won't equal 22.4 liters. But at 0 degrees celsius and 1 atm the mole of nitrogen is 22.4 liters. And a mole of xenon is 22.4 liters. And a mole of neon is 22.4 liters. And a mole of oxygen is 22.4 liters. And a mole of carbon dioxide is how much volume? 22.4 liters. And a mole of, I don't know, what's the next one? Argon is what? 22.4 liters. Yeah, it's pretty funny, right? 22.4 liters. It's hilarious. So hopefully everybody's got that. That'll be the last thing we do today. Study hard. Okay, I'll see you guys tonight from 6 to 8. So there's not going to be any office hours after class today. Remember those of you who are in my Monday lab, especially because I care about the most. So we have the lab practical today. So normal lab time, don't meet in the lecture hall, just meet in the lab room and we'll line you guys up, okay? Yeah, this one is just the two pool. Oh, no sign in Shido yesterday. That time maybe it'll be an easy question if you remember it, but do you think you'll put a question like that on the exam to kind of test the 22.4? Like if you remember that one mole at 273K, it is something that we've gone over in class, you know. Very cool, because as long as you remember it it'll be pretty easy to solve something like that.