 Let's go ahead and start today. The first thing I want to talk about is, well the first thing I guess we could talk about is the owl tube. I think I pushed it back till Friday night. So I think I'll give you guys enough time to finish all the stuff. It's not as bad as owl one as far as I can tell. So just let me know how that stuff is going. I think we're kind of working on that business. The blackboard side of the class 2, how we're going to take quiz 2. And for those of you who are listening to this right now, don't worry about it. I have to do this for the online people. For the in-class people, this is how to take quiz 2. So you're going to come to the computer lab HSC between these times. So you're going to have from Thursday, so tomorrow 8am to 10-15am. But the quiz is a half an hour long quiz. So if you want to do it tomorrow, you've got to get there by 9.45am. The thing is, is I have two classes taking a quiz in the same room over these four hours. So it's their class time tomorrow that we'll be taking up from this 8am to 9.15am. So they're going to have priority during that time. During your class time, which would be Friday 9-10, you guys will have priority. So if anybody comes during that time when you're supposed to be having class and there's not very many computers, you would get the computer before that. Otherwise, not during those two class times, not during 8-9.15 tomorrow and 9-10 the next day on Friday. It's first come, first serve. So become outside of those two times first come, first serve. But your time that you have priority is going to be Friday during lecture time. And it says that it's a 30-minute quiz covering the material from chapters 1 to 3 that we cover in lecture and through the album work. We haven't done very much of three in lecture. In fact, we haven't done anything. We might get some of it started today. And it'll just be that little bit that we get started. So it's not going to be any mass chapter three stuff. So it's going to be almost exclusively chapters one and two. You might get one or two questions. Make sure that you bring these required items. Photo ID, pencil. Because I don't know all of you guys yet. I'll name some of these as photo ID. A pencil calculator. And I'll bring blank scrap paper for you guys. And I'll also bring a periodic table for you guys to use that has the names of all the elements on it. You may or may not need them. So you guys got four hours to take a one half hour quiz. So make sure you get here before the half hour, before you have a half hour to spare. So don't come at like 10, 15 expecting it to take a half an hour because I only have that room reserved for so long. So are there any questions about that? No class on Friday? Go down and take the quiz. One half hour. But the questions, some of the questions will take you like a half a second to answer. Some of them may take you five minutes to answer. You're half a second or I don't have to answer? I think you could drop. Well, I mean, okay. So if it takes you five minutes to say ice melting to water, is that a chemical or a physical property? Then it may take you five minutes for those types of questions. But for me, and probably you, it will only take half a second. I think you know that type of stuff. So there's going to be a variety of questions on the test from these basic more straightforward concepts to some mathematical concepts. What I would recommend again is to bring a pencil and I'll provide you with some paper so you guys can write out any calculations. First, this is going to be through Al. So the typing in the answer is very important. Okay, that's the problem. You only have one shot at it. So we're going to do it. I think you guys are pretty good at typing in Al answers by now. We'll see how it works. I think it'll be okay though. So don't worry too much about it. This time we'll think of something a little bit different. Any questions? No questions, okay. Isotopes, the same element, but they have different masses. Just like chlorine has these two isotopes here, chlorine 35 and chlorine 37. Notice I'm saying chlorine 35 and chlorine 37. When I refer to an atom in such a fashion, the number that I'm referring to after that I state the name of the atom is referring to its mass number. So if I say chlorine 35, that means it has a total of 35 neutrons and protons in the nucleus. So it's the mass number. So if I say chlorine 37, I'm referring to the total number of subatomic particles in the nucleus. So of course chlorine 35 or 37 will always have 17 protons in its nucleus because if it didn't, it wouldn't be chlorine. But the difference of course in the mass number then has to be due to the fact that it has a different number of neutrons in its nucleus, okay chlorine? 35 of course we know has 17 protons so it must have 18 neutrons because 35 minus 17 is 18. And 37 minus 17 is 20 so chlorine 37 must have 20 neutrons in its nucleus. So you can see here, here's some more common isotopes. The isotopes of carbon. If anybody's ever heard of carbon dating, they use one of these isotopes to figure out how many years something has been dead for and so they can kind of estimate the age of certain things. Isotopes of hydrogen. Again this is used for some dating procedures because tritium hydrogen 3 here has only been around since atomic testing. So you can kind of radioactively date things in between the time atomic testing started to the present day. And by looking at the tritium concentration within these samples. This is due to radioactive decay of these particular isotopes that actually radioactively decay and then disappear over time. So you can tell by the concentration of something one of these isotopes in something when it was actually breathing it and being alive, especially the carbon 12. We're not going to go into very much depth about that right now. Maybe later in the term if we have some time we'll hit on radioactive dating and isotope decomposition. But for right now, just know that isotopes have the same atomic number, of course, but different mass numbers and they're at the same level. Okay, so we've been over this type of nomenclature before. Remember, Z down here is the atomic number. The mass number is up in the top left corner. The symbol of the atom is the big thing in the middle. And then the charge of the particle, which we haven't really gone over, but maybe if we get to it in next chapter we'll talk about what are known as ions. Ions. And they're actually atoms with a charge that's located in that position. So this is the way you would write it, like if you were writing it on a piece of paper or something. Again, if you wanted to write 4 on 11, a common way to write it would be just have the 11 in the top left corner be as the symbol of the element, not showing the atomic number because the atomic number and the elemental symbol of course contain redundant information. And notice 4 on 11 when we show it that way is uncharged because it doesn't have anything like that. And then you can know by the atomic number, the number of protons, the number of neutrons in the particular element. And notice here we've got two isotopes of 4 on 1, 4 on 11, 1, 4 on 12. Both of them have the different amounts of neutrons but the same amount of protons. So is everybody clear on isotopes and everything like that? Okay, so let's talk about relative masses. And I know everybody was very frustrated maybe with calculating, you know, the distance or the difference, I'm sorry, converting like miles to kilometers or certain values like that because on a surface they really don't have very much to do or very much in common with your thinking about what chemistry would be, okay? And we really start you to do those calculations with those types of figures because you're more familiar with those types of numbers and units than you are with the units that you actually need to use on a more common basis in chemical analysis, okay? But the thing is when you get used to doing those types of calculations that we did in Chapter 1 they really do help you along with the types of calculations that you'll be able to be doing or that you'll be required to be doing through chemical analysis, okay? So actually what happens is they give you kind of clues of how to do these new problems, okay? So the new problems are that you do them exactly the same as the old problems they just have different units, okay? And those units might be a little more confusing because we're not used to them as well, okay? The first one we're going to talk about one of those units is the mole, okay? So let's talk about relative masses first of all and learn why we need to come up with a new unit called the mole to help us describe the way that atoms are counted, okay? So the extremely small size of atoms and molecules so of course we can't see them there so there are unlike things that we're normally used to dealing with on an everyday basis, okay? So they're very small. So it makes them inconvenient to use their actual masses for measurement. So of course their actual masses are on the order of 10 to the negative 24th and 10 to the negative 26 grams. So those are extremely, extremely small numbers. In fact, we don't have really any relative basis on how to measure that or to ask ourselves, you know, well, how big really is that, okay? So we really need to come up with a different sort of thinking about it so we can start to say, okay, yeah, I kind of am understanding how the magnitude of the mass of these types of things, okay? So, yeah, so it makes it inconvenient to use their actual masses for these measurements or calculations so we use relative masses. So we're actually going to compare the atoms to each other, okay? And we're going to come up with a new unit called the AMU or you might see it as U, okay? It's the same thing. So relative masses are a comparison of the actual masses to each other. For example, if an object had twice the mass of another object, the relative masses would be 2 to 1, okay? Does that make sense? 2 to 1? If something is twice as heavy or as something else. So the relative mass scale of the unit that we use is this atomic mass unit or the AMU and it's defined as one-twelfth of the mass of a carbon-12 atom, okay? So a carbon-12 atom is exactly 12.000 AMU to infinity, 00 whatever AMU. So one-twelfth of that would be one AMU. So that's the official definition of one AMU. But one AMU can be thought of to represent the mass of either one proton or one neutron, okay? So hydrogen one has one proton and its mass is one AMU. Helium two would have, or helium four would have two protons and two neutrons, the mass being 4.000 AMU, okay? And then, so if you could look, yeah, again, an atom with the mass equal to twice the mass of carbon would have a relative mass of 24 AMU. So that would be like, and then there's this conversion factor here that really converts AMUs to grams. So this is something that you should try to remember, but I'll give you this conversion on any sort of text, okay? I mean, I want you to understand more the concept instead of memorizing all these numbers and conversions, okay? But it is good to kind of remember this number because it gives you an actual idea of how small these things actually are. Okay, so I'll give you this calculation, but let's try to use this calculation in a problem, okay? So let's try to figure out how many grams, I don't know, one of Vandium 51 atom weight. So Vandium 51, like that, let's draw the whole, okay? So if I were to ask you what's the mass of Vandium 51 atom in AMUs, you would tell me what? Who would tell me what's the right answer? Vandium 51 AMUs, right? So this has got 23 protons, right? And then it would be 28 neutrons, right? 23 protons, 51, okay? And remember we said this mass number is the total number of protons and neutrons and each proton and each neutron weighs how many AMUs? Each. One, right? One AMU each. So the total mass of a Vandium 51 atom in AMUs would be 51 AMUs. Does that make sense? 51. We even go so far as to say per one Vandium 51 if we wanted to go so far, okay? And use that as another unit. Well, how many grams does one Vandium atom weigh? So how will we do that? Well, we just use our conversion factor that we have up there and say, well, multiply that by, well, one AMU equals... Does everybody see how I did that? It makes sense, right? So then AMUs, of course, will cancel out and now we have units of grams per one Vandium 51, and that's what we were really looking for, right? Figures of the AMUs here, okay? Because this is a conversion factor. So it's going to be 8.5 into 23 grams if you want to. Very small amount, very small. I can't even imagine. Like relative to this piece of chalk, it's like nothing, you know? And this piece of chalk in my hand is like nothing, you know? So you can imagine. I don't know how familiar you guys are with holy pieces of chalk, but it's not very good. So what if... I mean, how many Vandium atoms do we have to give to have one gram in our hand, which is even less than probably this piece of chalk? It could be a lot, right? A lot of Vandium atoms. And in fact, some other guy... I might be jumping ahead here, but some other guy figured it out. His name was Avogadro. He used Avogadro's number. Has everybody heard of Avogadro's number? He used this number to figure out, to say, well, you know, I can figure out how many atoms we need to give the exact atomic weight of each of those atoms. And it's always the same number, okay? It's actually 6.022 times 10 to the 23rd. That's a very large amount of atoms, okay? So Avogadro's number, 6.0. This number, in fact, doesn't have to be atoms. It could be anything. It could be people. It could be pieces of chalk. It could be anything, okay? So this is just another unit. This would be...this is now known as the mole. Okay? One mole is 6.022 times 10 to the 23rd, whatever. Okay? Like, donuts would be a mole of donuts, okay? Or atoms would be a mole of atoms. Just like if we had a dozen atoms, it would be 12 atoms, okay? So this is just a number, just like a dozen is a number. Or a baker's dozen is a different number. Or a pear is a different number, okay? Or, I don't know, if you guys can think of another number type unit, right? So one mole, mole is just a number. And that number is 6.022 times 10 to the 23rd. Okay? In this case, we're going to look at it as a mole of vanadium 51 atoms. Let's do that. Okay? And we'll see how much the mole of vanadium 51 atoms actually weighs. And we'll see if it weighs something that we can finally get our head around. Some number that we can get our head around. So we know that one vanadium 51 atom weighs 8.5 times 10 to the negative 23rd grams. So we'll just use that. And this inversion factor to figure out, well, how many grams does one mole of vanadium weigh? Okay? So just use that number, 8.5. If you want to know grams per mole, we want to cancel out vanadium atoms, right? So we say 6.0. We're going to cancel with that. Everybody see that? Canceling. And now we've got how many grams per mole of these atoms we've got? 51.101, you know? So it all depends on what numbers you rounded as before. But it doesn't matter. The main point, because we're going to only take it to two significant digits, right? It's going to be this 51. And notice, so it'll be 51 grams per one mole of vanadium atoms. Notice what's the similarity between this number and the original number up there, 51 AMU. They're the same number, right? But they have different units. And in fact, what Avogadro figured out is that if you multiply the mass by this number of atoms, you'll always get the atomic mass in the unit's grams per mole. So you can kind of do a shortcut here. And if you ask, well, how many grams does one mole of mercury weigh, you can say 259. You can just write that number off of the periodic table. I suggest that you go through a few of these calculations on your own. But this is what it means to use the unit's mole. Mole is just like any other unit, like a dozen or a pair or something like that. It's just describing a number. It's not describing anything else in that. But it is a useful number because it allows you to say the mass number of the atom that I'm looking at is equal. We can set it equal to the molecular weight, which is grams per mole. And we can use this. This is much more useful of a term to use than this is up here because we measure in the lab things on the gram scale and not on the AMU scale. Does that make sense? So it kind of helps us say, if I weigh 83 grams of vanadium, how many actual atoms do I have? So I don't have to go back and say per 51 AMU. So it really does help you. Right now it might seem like a lot of excess work, but in the end it will help you a lot. So make sure you know how to do this. I kind of skipped probably a way ahead and did this whole problem, but let's just go over all that stuff that we just talked about. So atomic weight, that's the mass or the number that you see at the bottom of each of these symbols on the periodic table. That's the atomic weight. And that's the weighted average of the masses of all the isotopes of an element based on natural abundance. I'm not going to go too heavy into figuring out the average atomic weight. I think I put maybe some al-problems, but on the test don't worry too much about it. Okay? Not figuring it out. But knowing what it is is very important. Okay? And knowing what it means. So weighted average is the average corrected by the relative amount of each isotope. So chlorine, of course we know has two isotopes. One is chlorine 35, one is chlorine 37. And notice the average atomic weight of chlorine is 35.452 cents. So it's somewhere in between 35 and 37. On the side of 35, because there's more chlorine 35 than there is chlorine 37. Again, look at this here. And you can figure it out. It happens in three to one ratio. You can figure it out by doing this calculation here. By converting the percentages that you get to a decimal and then multiplying by the relative masses. Okay? So if you look at chlorine 35 here, this percentage is 75.77%. So if you divide that by 100%, you're converting that to a decimal. Then all you got to do is multiply by its mass in AMU. And when you do that, you get this number. You do it to the other one as well. And then you add these two numbers together. And that gives you the weighted average. And that gives you that number that you find at the bottom of the periodic table. So you can go through that one. There's another one here. What's the atomic mass of boron? And then you can try it on your own. Or the atomic mass is the weighted average of these things. So you would convert this to a decimal. Convert this to a decimal. Multiply it by, in this case, 14. In this case, by 15. Add those two together. And you should get somewhere around that number. Notice somewhere around, it says 99.63% of 9 to the 14. And look how close that number is to 14. It's very, very close. Because there's only a tiny bit of 9 to the 15. Okay? So that tiny bit of 9 to the 15 actually raises the average atomic weight by .0067 volt. If you could try this on your own. If you've got any problems, ask me after class and we can do it together. Of course, it's the mass of a molecule. You can find the molecular weight by knowing the molecular structure of the molecule. So, for example, one that we're very familiar with is water, which is H2O. Okay? So we can figure out what the total mass of a water molecule is by adding the average atomic weight of the atoms that it's composed of. Okay? Just like it says here. The molecular weight of water is 2 times H because you've got two H's in it. And plus 1 times O, which is because you've got one O in it. Okay? And then if you add those two together, you get, well, the relative weight of hydrogen is 1.0018. Okay? A lot of people put, I'm sorry, 1.008. A lot of people put. And the relative weight of oxygen is 16. You'll find that a lot on periodic tables. So if you combine these together, you'll get 18.02, or the molecular weight of water. Does that make sense on how to figure out molecular weight? Okay. So I guess let's try it for carbon dioxide sulfide together. And let's, I want to go back to this one. Let's try this one together, too. Okay? This is nitrogen. One. Does anybody get the answer to this? And do you get it close to that number of atoms? Let's try it together. And then we'll try that molecular weight together. The first thing you want to do, remember, is to convert that to a decimal. So we take 99, what is it? 99.63%. And then the way to convert that to a decimal is you divide it by 100%. Because the percent is the unit. Okay? So we've got to get that out. So when we do that, we get 0.9963. Okay? So no units, notice. And then we'll do the same thing to the other one, which is 0.37%. We've got to divide that by 100. Okay? Because the units cancel here. And then we multiply this by, for the nitrogen 14, we're going to multiply that by 14.00. Remember, if it's 14, the O's are infinite. So you can take it to as many significant digits as you want, really. If you don't have these numbers over here. So this one you're going to have to take to four significant digits, this one to two. Okay? So watch that. But this one is an exact number. So you don't have to worry about the number of significant digits there. So we'll take that, multiply that by 14 AMU. Multiply this one by 15 AMU. And this one down here, just to figure it out eventually. We'll take this. We'll just add these two numbers together. And I use the original numbers. I didn't reduce them to these number of significant figures. But it doesn't matter, because eventually you'll get about the same thing. 14.0037 AMU for the average mass. But unfortunately, this is going to have to go to two significant figures. So two significant figures to really figure it out. But if we had taken this one, maybe to more significant figures, we could actually see that it was doing something. But you can really, it's evident, looking at this portion. Let's go ahead and figure out the molecular weight of carbon disulfide. Okay, so carbon disulfide is colorless liquid CS2. So what's the atomic weight of carbon? The average atomic weight is what we'll use. 12.01, yeah. It's always going to be dependent on the particular periodic table that he used, because they all estimated differently. But carbon is 12.01 AMU. Carbon disulfide, molecular formula CS2. You don't know how to figure that out. Yeah, eventually you'll be able to figure that out. And all we've got to do is say, okay, one times 12.01 AMU, you know, 6 AMU. Probably, we've already done it. What is it? 6.13. Everybody understand how to do that. Okay, so this is what essentially we did with vanadium 51 quite some time ago. Okay, this time is, instead we're doing it with chlorine 19. Remember, the first thing we did was ask, well, how, or what's the amount of, what's the mass in grams of one vanadium 51 atom? Well, here we've got a calculation that's similar, but instead we're asking what's the mass in grams of one chlorine 19. Okay, so notice here we've got chlorine 19 multiplied by that conversion factor that we gave you. What happens, AMU of F cancel out, and we get grams per one chlorine atom. Okay? Of course, we don't only measure one chlorine atom when we're working in the laboratories. Almost impossible, okay? And last year, you know, very, very specialized person. Okay, but in the general chemistry lab, we're never going to be just looking at one atom, okay? We're going to be looking at number of atoms, in fact, a very big number of atoms, because one gram, as you saw, was quite a big number. Okay? So, again, we're going to use this new term known as Apagodra's number, which is 6.022 times 10 to the 23rd. That is a number you will have to know, okay? I won't give that to you on a test or anything like that. Different than that conversion factor, from grams to AMUs. That I will give you, but Apagodra's number is one of those things that, I don't know, once you learn it in chemistry class, you should take to your grade, or whatever. Okay? So, everybody should know that for as long as they live, as far as I'm feeling. Okay? So, this new unit is needed to deal with the collection of items that have such large numbers. But it's just like saying a dozen of something. It's just a number, okay? So, I can have a mole of, you know, I don't know, cans of Coke or something like that if I wanted to. And that would be 6.022 times 10 to the 23rd cans of Coke. Okay? Or donuts. Or whatever you want. And remember, when we did that Vandium 51 calculation, when we multiplied it the number of grams per Vandium atom times that conversion factor, what did we get? We got 51 grams per one mole of Vandium. Remember that? So, it always comes out to the same number that you find on the periodic table as the mass number. Okay? So, if you will. Remember, if we look up here at Sodium, its average atomic mass is about 22.99. Does everybody see that? And if we multiply that number by 6.0 or the mass of one atom of Sodium by 6.022 times 10 to the 23rd will always get 22.99 grams per mole of Sodium. Okay? I know it seems obvious right now. I would recommend that you guys go through this at least a couple of calculation times, okay? And then make sure you kind of embody this sort of concept, okay? And then it will really help you guys out a lot more when you're working in the lab and kind of understanding, you know, the actual scale that we're working on. And you can see, of course, a mole of calcium would be 40.08 grams. Look at the average mass there. And a mole of sulfur, 30.07. Again, this is that similar calculation that we did with Fandium. And here's a cool picture I'd like to look at. This is one mole of each of these metals. So that's one mole of copper there. That's one mole of aluminum, okay? One mole of lead shot there. One mole of sulfur. One mole of chromium. One mole of magnesium. Okay, so those are all kind of amounts that we would, might be weighing in the lab or we could say maybe half of that amount or something like that. That might be an appropriate amount of material to be working with in the lab as opposed to one atom, you see. Okay, so it makes it very convenient because this is like right on our kind of hand scale. You know, it's not like a big dump truck of stuff and it's not, you know, microscopic amounts. Okay, it's about our hand scale. Okay, does that make sense? So this is why we use this number. Not to mention the fact that it comes out to be its atomic weight. And then you can apply that same concept to compounds. So you say, okay, well one mole of water equals 18.02 grams of water. 18.02 amu remember was the molecular weight of water. Okay, so it again is 6.02 times 10 to the 23rd molecules and that equals the molecular weight. Okay, so it's the same concept over and over and over. Okay, so get familiar with it. It really will help you out immensely learn this earlier. There are some more mole quantities. You can see a mole of water is 18, about 18 mils of water. Thanks to not very much, maybe a small swallow of water. Thanks. And then, so that's a molecule, right? That's not an atom. So you can apply that same concept. A mole of sodium chloride is this beaker. Very large amount. Beaker of aspirin here. That's one mole of aspirin. And nickel chloride hexahydrate. That's one mole of that. And you can see, you can imagine, well what would be the molecular weight of aspirin? Well it must be 180.2 amu. It must be because one mole of it weighs 180.2 grams. So embody this concept. I know I've been saying that over and over and over and I promise you will help me. And then, you can actually do on your own some of these calculations that we did earlier. Grams, the moles, the atoms. And then there's some practice calculations where you guys can do. And some more practice calculations. I think we went over how to do every one of these types of things in this lecture today. Okay? So, yeah, I guess we'll leave a minute or two early today. Okay? And if there's any questions, feel free to come up here.