 Let's just continue where we finished up last time. So like I said, I'll post all that information about your plan to online after class. But let's finish what we were talking about. I think the topic we were on was limiting reagent. So of course, the last thing we talked about was balancing equations, calculating moles to the moles, calculating grams to moles, and calculating grams to grams, OK? So hopefully you guys have practiced some of these types of problems using the chemical equation. Remember, you've got to balance your chemical equation first. And then do the mole to mole ratio thing. OK, recall we were talking about limiting reagent. It's the thing that gets used up fastest. So you can see here it's going to be a one-to-one mixture of hydrogen to fluorine. That's what the reaction says, a one mole for one mole to make HF of one mole. So let's just write that down just to make sure. If your total of the chemical equation indicates that we need one mole of hydrogen or one molecule of hydrogen, whatever way you want to look at it, for every mole or molecule of fluorine to make one mole of hydrochloric acid. But you see here we've got one, two, three, four, five, six moles of hydrogen and only four moles of fluorine. So we've got excess hydrogen, if you will. There's extra hydrogen molecules, OK? So when we combine these four with these six, of course, the four can only combine with four of these hydrogen molecules. So there's your mixture combining, right? And then we go over here. We can see we've got one, two, three, four, five, six, seven, eight HF molecules. But we still have some unreactive H2 molecules, OK? And that would be your excess reagent, OK? So this would be the excess excess reagent or reacting. And this would be, of course, your limiting reagent. Because when we came over here, we still had plus excess. Notice this isn't one of the products of the reaction. It just didn't get reacted. So here's a pretty interesting and convoluted problem. Calculate the maximum amount of SO2 that could be produced by reacting 55.2 grams of oxygen and was 50.8 grams of hydrogen sulfide. So I haven't tried this problem. The first thing you're going to have to do in this problem, so this one's a little more difficult than I think I would give you on the test. I'd probably at least give you all of the components of the reaction equation. But anyway, so let's write this down. H2 left over. So to one to one ratio, right? So this is a balanced equation. That's the first thing you would have to do. The next thing you would have to do is, well, maximum amount, this term amount, this usually refers to moles. So that's what we're looking for here, is the number of moles of SO2. OK, well, we know a couple of things. We know that we have 55.2 grams of oxygen, so the mass of oxygen. Remember, this is the moles to moles to moles to moles ratio. OK, so mass doesn't help us out there. We need to convert it to moles. OK, so the first thing we want to do is convert this to moles. How do we do that? Well, we look up there on the periodic table. The periodic table gives us the atomic weight of oxygen. Of course, we multiply that by 2 to get O2. So it's going to be 16 times 2 or 32. I remember, so it's 1 mole of O2 equals 16.0 grams of O2. I mean, 32, yeah. OK, so we know from this equation, this will give us, now, we don't have the mass anymore. We've got the number of moles. So now we need to figure out, well, what's the number of moles? So why are we doing this? Why do we have to figure out both of them? Because we need to know which one's the limiting reagent. Which one's going to be present in the least amount. So what do we have? We've got 50, 50.8. Which one of these two is the limiting reagent? So remember, what is the limiting reagent? The one that's in smaller amounts. So since this is a 1 to 1 ratio, right? 1 to 1 ratio, then we can compare these things directly. So the one that's in the smaller amount is that one. So that's the limiting reagent. OK, so we're almost there to figure out the maximum amount of SO2 that can be formed. Well, the maximum amount of SO2, since we got what, the limiting reagent is H2S, since there's a 1 to 1 ratio of H2S to SO2, all we've got to do to figure out this is say, well, the maximum amount of reactant that could have reacted is 1.490 moles H2S. The origin factor that we got by the reaction equation, 1 mole of H2S, holds 1 mole by 1. Figure out which one was the limiting reagent. This reaction equation was different than this. If it was like a 2 to 1 ratio of these guys, we would have to factor that in here. So if we needed 2 H2S for every O2 here, we would have to multiply this by 2 to indicate which one's the limiting reagent. And in that case, the other one would be. So watch that when you've got different numbers of or different coefficients in your reaction equation besides just 1. So I think the last thing we're going to talk about and this chapter about calculations involving the chemical equation is reaction yields. The theoretical yield for any reaction. So I kind of want to keep this stuff up here. But the theoretical yield for any reaction is going to be the number of moles that the limiting reagent gives you. So that would be like 100% of your reactants when the products. So in this case, the reason I want to keep this up here is because it really does help to indicate what we're talking about here. So in this case, the theoretical amount that could be made is 1.4090 moles. So that would be theoretical. The actual yield is spelled wrong there too. The actual yield is going to be what you get in the lab. So if I did this reaction and I didn't get 1.4090 moles of SO2, but this is just the number I'm coming up with out of my head. Let's pretend instead I got 1.001 moles. So this is what I got. It could be a problem, right? I would have to give you this number. So let's go back and figure it out. Remember, the theoretical is this, and that's the actual guarantee there'll be a couple. All we do is just plug in these numbers. So the actual, the theoretical is now the new units of this, 1.001 divided by 1.4090. So if you're getting over 100%, you're probably not doing the problem, right? Because of course, 100% is the maximum amount you can get in any reaction. Why is that? Because of the law of conservation of mass, right? You cannot create or destroy matter. So if you're getting over 100%, you've done something wrong. So does everybody see how to do this type of a problem? Hopefully you can do that. Remember, this part of the information here, I'd have to give you in the problem, OK? So what happened to the rest of this? So we're missing like 33% almost here, right? What happened to that 33%? Well, it could be either unreacted starting material. So it could have stayed in the reaction mixture as unreacted starting material. Or reactants don't always just go to one product, OK? They can go to what we call side products, OK? And in this case, probably a lot of that occurred, if this was our yield. Or we lost some of it, so we dropped it or whatever. OK, here's another calculation for you guys to try. You can do it not only with mass, but with the moles. But if you have the masses, you can compare them directly, OK? But of course, they have to be the mass of the same component of the reaction. Because if they're different components, if you've got the mass of Al and Fe, those are way differently. And their relative weight, whether you're out of table, is different. So you're going to have to convert them to moles. But if you've got mass and mass of the same product, then you can do this as well. So you could go over that one on your own and then try this one on your own. So there's terms that we use. If we went to 100% and we got all of the reactants to go to products, which of course that never happens in an actual reaction, we normally get some loss of products from reactants. So not all of the limiting reagent, if you will, went to products. But if this is the case, so this theoretical type reaction condition, we call this a quantitative reaction. If all reactants go to products. That's a question. Uh-huh. OK. Because that doesn't account, that discounts the relative weight factor, OK? But if you've got the same product, right? If it says you were expecting 30 grams of aluminum and you got 25 grams of aluminum, you don't have to convert that to moles, OK? So that's what I was saying. So if you've got the same reactant or product that gives you both the actual and the theoretical yield for in grams or in mass, you can use that mass. But if it's like two different ones, right? Then you can't use that. You're going to have to convert. So like in this equation here, you have the one that we did, the big one on the board. You have to convert that to moles, OK? It's best if you convert everything to moles. But I know on a test, you're like going as fast as you can. You're like trying to cut down as many steps as possible. So if it gives you like that, then don't convert it, you know, because you want to cancel that step out, OK? OK, anyway, so going back to this, quantitative reaction, all products go to products. Non-quantitative reaction, some of the reactions are left over, OK? So what does that mean? It's either that, you know, you didn't get, you know, full what you were expecting, like 67.8%. So it could be some of the reactants are left over, or some of those reactants could have gone into side reactions, OK? So it could be non-quantitative either way. So if all your reactants get used up, right, but they went to side reactions, it's still non-quantitative. So the theoretical yield, again, is the maximum amount of product, and the actual yield is something that you perform in the laboratory and I have to give you on a problem, OK? So go over those extra ones on your own this weekend. Chapter 6, states of matter, yeah, they're out today. So the states of matter, Chapter 6. So if you recall from Chapter 1, so again, we're going back to Chapter 1, thinking about Chapter 1. So this class is always, you know, going back and forth because you need to know that stuff from Chapter 1 to do stuff now. But remember, in Chapter 1, we just describe changes in state. Remember, state is like phase, right? So you can change your state by going from a gas to a liquid or a liquid to a solid or a solid to a liquid or a solid to a gas or a gas to a salt, OK? So these changes in state, they're physical changes. Here we have pictures of a solid, gas, and like, so this is what we're going to be talking about. And really talking about them, kind of what we were talking about, like how we were talking about in the experiment from a week ago, with thinking about how molecules move, OK? So during a change of physical state, many other physical properties also may change. Density, shape, compressibility, and thermal expansion. Of course, hopefully, you know the density could change, right? You know these things, OK? So you know that the density of water is less than the density of ice, right? Because ice floats on water, OK? So there's a difference in the physical properties among the states that you guys are already aware of what we're going to beat it to death. So remember, density is mass over volume, OK? So if you can imagine the mass of these things in the beakers, these particles in the beakers. So say this is just the same amount of particles going from a solid to a liquid to a gas. So if we imagine that the mass of these three beakers or these three Erlenmeyer flask is the same, right? The stuff that's in the flask is the same. We can see that the volume has expanded in all of it, right? So if we say that the mass of this is 1, the mass of this is 1, and the mass of this is 1, right? The volume of this is very small. Volume of this is bigger. The volume of this is bigger, right? So if we take the same mass and divide it by increasingly bigger volumes, what we'll get is a less and less density, OK? So what we find here is that we've got very dense solid, OK? Not as dense liquid and very not dense gas, OK? Notice in water, the solid and the liquid state are opposite in density, OK? So that's a weird one. Usually you'll have the solid being more dense than the liquid. With water, it's different. That's why, again, the water or the ice molecule or the ice particles flow on top of the water molecules. OK, so hopefully, looking at this, you can see the shape has also changed, right? Going from just kind of a, you know, regular solid shape to more disordered state in a liquid, kind of a free-flowing state. And then from that to the shape of the Euler-Meyer flask. The shape of the matter takes, depends upon the physical state of that matter. The compressibility, this is the ability to squish it down, right? Compress it on top of itself. The change in volume of a sample that results from the change in pressure, the compressibility of a gas is much greater than that of a liquid and then a solid is very little able to be compressed because it's due to the space between the actual particles. With a gas, if we look, the space is very large between the different particles, right? So you can imagine being able to squish those things upon themselves, OK? A liquid, you know, it's got a little bit of space that we can kind of push it down, but not very much, honestly. And in a solid, they're all stuck together already. So it's almost impossible to compress a solid, especially relative to, you know, a liquid or a gas. And then thermal expansion, that's exactly what's happening here. As we increase the temperature or the average kinetic energy of these molecules, they go from this state to this state to this state. So that's the thermal, thermal meaning heat, right? Expansion meaning getting bigger, right? So these things thermally expand, right? Going from a solid to a liquid and then from a liquid to a gas. So here's kind of a table that goes over all of that stuff that we just went over. So it kind of puts it all right in front of you at once. And here's another table that talks about a couple of other things, the particle motion. So the particle motion for a solid is that they vibrate around a fixed position for a liquid. The atoms remember roll over or slide past each other. And particle motion and a gas, they don't even feel each other. They're essentially very far apart from each other. The intermolecular distance is very large. So they don't even feel any sort of attraction or repulsion from each other. For a liquid, the molecules are close together. And for a solid, of course, they're also close together. Very close together. And here you can see the changes in state from a macromolecular point, like we've been looking at for the last few slides. So what we actually see in an everyday experience. So you see the solid has the fixed shape, fixed volume. Liquid has fixed volume, shape of the container. Gas, volume of the container, shape of the container. But we also can see it here on the molecular level, where the solid is very structured and rigid. Okay, the liquid is less so structured, but the molecules are still very close together and kind of touching each other. And then you see when it goes to gas, the molecules are far apart from each other. They don't even interact with each other. So this is a different view, right? So this is what we see with our eyes. And this is what's actually happening on the molecular level. So why does this stuff happen? Well, there's this theory, of course, that's been, you know, worked on over the last couple of centuries. It's known as the kinetic molecular theory. And this describes actually what particles are doing in the various phases, okay? So the kinetic molecular theory of matter is a useful tool for explaining observed properties of matter. So the first postulate is that matter is made up of tiny particles called molecules. So hopefully by now you believe that. Particles are matter that are in constant motion and therefore possess kinetic energy. So anything that's in motion, anything that's in motion possesses this thing, this attribute or whatever called kinetic energy. Everything in motion, yeah, everything in motion. Kinetic energy is just, again, it's a measure of, you can think of relatively the temperature of these molecules, okay, or the amount they're vibrating. Postulate three says that the particles possess potential energy as a result of repelling or attracting each other. So the potential energy is the energy that these particles actually have stored within them. So if I have, if I'm a negative particle, I have the potential to attract a positive particle, okay? That's what potential energy means, is that you've got the potential to do this, even though you're not doing it at this point, okay? Postulate four, the average particle speed increases as the temperature increases. Remember the temperature is again kind of a measure of the average kinetic energy. And particles transfer energy to one another during collisions in which no net energy is gained or lost from the system. So they're like these pool balls that are colliding with each other but not losing any energy on collision, okay? So they're like on a frictionless pool table hitting each other and when they hit each other, they don't lose any energy either. So they just keep bouncing and bouncing and bouncing. That would be like if they were in a container, okay? So let's talk more about kinetic energy. Again, kinetic energy is the energy of a particle that will pass as the result of being in motion. It's usually abbreviated by this capital K E and it's calculated using this equation here. K E equals the mass of the particle times its velocity squared divided by 2. What you need to know about kinetic energy or energy in general is that it's usually given in the units of joules. So you're going to have to do some unit conversions here. Notice one joule equals what we know as one, this capital J is joule, it's spelled like this, some guy, Dr. Joule. And that equals one Newton, Newton, that's that capital N. That's also named after, you know, the long-sense dead guy. Dr. Newton or Isaac Newton, if you guys have heard of that guy. So Newton's times meters are Newton meters. One kilogram meter squared over second squared. So what you're going to be doing, what you're going to be getting in this equation here, notice, it's going to be kilograms, that's the mass unit, right? Kilograms is the mass unit. Velocity, that's meters per second, okay? So when you square that, meters per second, you're going to get this meter squared second squared. So after you do this equation here, you'll get these units here, one kilogram times meter squared over second squared, and you'll want to convert that to joules, okay? And it's a one to one ratio. So the mass is given in grams, you're going to have to convert it to kilograms, okay? Or if the time is given in minutes, you'll have to convert that to seconds, okay? So watch out for that. Okay, so let's try one here, I'm giving you the answer there, but it's just a matter right now of plugging these values into this equation. Look, we've got this equation, kinetic energy times the velocity squared over two, okay? And it says the average kinetic energy, what is that, of an electron, well, an electron mass. So how do we plug those things into this to get that number? All we do is take what it says mass, put the mass of the electron. So I'd like you guys to try that on your calculators. You should have your calculators out anyways. Try that on your calculators and see if you get that number. Okay, notice that the units that you're getting here, right, if I multiply kilograms over seconds squared, kilograms, second, yeah, because it's meters per second. And that's why I wanted to do this to make sure you guys know. When it says this, right, when it says something like this, it means that, right, divided by this. Divided the whole thing by two. So the first thing I did was this, plug this into my calculator squared, right, multiply that by this, okay, then divided the whole thing by two, then I got these units out of that equation. But then I went back and remembered that if I got kilogram meters squared over second squared, that equals joules, okay. Again, if I were to have given you this in grams, right, you would have to convert that to kilograms. Or if I were to have given you this in minutes or whatever, miles or whatever, you know. So it has to be these units. Or you won't get joules out of them. Okay, so try this one on your own. It's a little more difficult. What's the velocity of a hydrogen molecule? I'll give you the mass of it. If the kinetic energy of that molecule is 3.683 times 10 to the negative 24th joules. Try that one on your own Monday, if you like, okay. Okay, so that's enough about kinetic energy and that equation. Let's talk about potential energy again and these cohesive and disruptive forces. So remember the potential energy that the energy of particle has as the result of being attracted or repelled by other particles. So if I'm positively charged, I have the potential to attract negatively charged particles, okay. That's what that's saying. So a cohesive force, we use these terms cohesion and disruption. Cohesive force is known as an attractive force between molecules, as you would expect, okay. This is associated with potential energy. A disruptive force results from particle motion. It's associated with kinetic energy. Okay, that's like particles trying to get away from each other. Okay, because you've raised their average kinetic energy. So if you recall, we've talked about different types of forces already. We're going to review that a little bit. We've talked about bonds, bonds. So that would be like in the water molecule, you've got two bonds between the oxygen and the hydrogen. How to build water on other small particles. So the intramolecular forces are the molecules, are the forces that exist within the molecule, that holds it together, right. So this whole molecule is one unit. The things that hold these atoms together are these things here, these lines, they're intramolecular forces, we call those bonds. It's specifically covalent bonds. So these are the forces responsible for the chemical properties of things. So the reactivity is, remember, when you react you've got to break a bond and make a bond. Okay, specifically a covalent bond. Okay, intermolecular forces on the other hand, we talked about these before already, specifically hydrogen bonding, but that's what we'll look at right now. Because when I have two molecules, the attractive forces between those two molecules. Okay, these are known as intermolecular forces. And in fact there are a wide variety of intermolecular forces. You can see one of them up on the board, that one's called hydrogen bonding. But the different intermolecular forces, the result of those is the actual physical properties of these substances, okay. So in tromolecular or bonds, that has to do with the reactivity of these molecules. The intermolecular forces has to do with the changes of physical state or the physical nature of these molecules. Okay, so make sure you get those straight. Okay, so a phase we've talked about is a physically distinctive form or state of matter of a substance. You can see in this picture, all three phases of water are present, solid, liquid, and gas. A phase change is a physical change from one state to another. Here are the phase changes, the names of them, condensation, vaporization, freezing, melting. Melting has another name used. Okay, probably most of you are familiar with this, but you can see what we call a dynamic equilibrium. Okay, the dynamic equilibrium means it goes back and forth, back and forth, back and forth. Okay, so you can see that ability of having, establishing the dynamic equilibrium between the gas and the liquid phase or the solid and the liquid phase, or there are subsubstitutes that can go back from solid to gas, solid to gas like carbon dioxide, okay, or dry ice. Okay, so let's look at the definitions of some, we'll go over dynamic equilibrium too, but. So remember, a gas, according to the kinetic molecular view, a gas has the energy of attraction or the cohesive forces that are extremely small relative to the particle speed. So the potential energy is small, kinetic energy is small, kinetic energy is small, potential energy is small, kinetic energy is high, okay. Liquid, the energy of attraction or the cohesive force is stronger because the particles are in contact with each other, but they still have enough kinetic energy to roll and tumble over each other. So the kinetic energy and the potential energy are more closely related and illiquid, okay. And then if we look at the solid, the energy of attraction, the cohesive force is extremely high. So the potential energy is really high, the kinetic energy is really low because they don't move very fast, okay. So that's the difference between potential and kinetic energy. And then let's look at this process of the dynamic equilibrium. You can see before, when you just put the liquid in there, if you've evacuated this chamber, in case of no air or no nothing is in there, pour your liquid in there, you can see at the beginning, first molecules from the liquid will evaporate to form a gas layer. Over time, once that gas layer is established, of course the volume of the liquid decreases slightly because the particles of the liquid have gone into the gas phase. But you see, remember we call this the dynamic equilibrium. The gas particles now going back into the liquid phase and the liquid particles going back into the gas phase at equal rates, okay. So dynamic means ever changing or moving around, okay. Equilibrium means an established going back and forth between, in this case, two physical states, okay. So we'll stop there for today. We'll go over some more of chapter six on Monday.