 Okay, guys. Good morning. There's a sign-in sheet going around already. And today is the big, but it hopefully knows what their score is by now. If not, they're posted on Blackboard, okay? The average for the exam was not bad. So for the in-class people was a 62.5 and for the online people was a 63. That's about where you think the average is going to be. There were four people that got hundreds on the exam. And so it wasn't impossible, right? This is like the sixth time I've taught this class and it's nothing different than I've given anybody else previously. So I can guarantee you that from previous classes also it wasn't impossible. There were many, many people that got A's and many, many people that got B's, many more. And many, many got C's too. So if you think it's impossible already, I can assure you that it's not, okay? The other thing I want to say is recall that my policy is just for this eventuality because most people at this point have underestimated the class before they took exam one and now they want to get their scores back to being higher. My policy is at the final if you get a better grade than any of your previous three exams you can replace that grade with the final grade, okay? So if you got like a 10 or 20 or something on this exam and you get 100 on the final then you get 100 on exam one, okay? So you're not in any dire straits, okay? So everybody should have their exam grade already. I posted the online people in Blackmore but and for those of you who are here of course everybody's still enrolled in the class. You can see how many people actually care about the class. So this is probably what's going to happen. You're going to see your grades go up. The problem is that a lot of people have now dropped the class so the average is also going to go up, okay? So the average and the high scores are what I set my curve by. So again you're just going to have to study harder if you feel you got a grade that you didn't want, okay? But if you're happy with the grade that you got then just continue studying the way and the amount that you've been studying. You can probably be honest with yourself, ask yourself how much you have studied and go back to that first quiz and see if you've lied to yourself or to me, okay? Anyways, and if you have been studying 10 hours a week and still aren't getting it then I suggest you just start studying more, okay? Because this is not an easy class. Anyways, the keys for the exams are here and here. So I gave two different exams. There are plenty of blank copies of the two exams. If you want to take one of those home with you and try it on your own some more. Of course the keys are also online, okay? So again the average is about a 63 for both classes. It's a pretty good average. There are a lot of low scores and a lot of high scores. So there was kind of two pumps of people. People that got 100s, 90s and high 80s and people that got 40s, 30s and 50s, okay? So very little in between. If you're in the latter section you should probably get in here. Anyways, I'll give the exams back about 10 minutes left in class, okay? So anyways, let's start talking about Chapter 5, Chemical Reaction. Okay, so the first thing we're going to start talking about is the chemical equation. A chemical equation is a shorthand notation of a chemical reaction. I could describe to you a chemical reaction in words like I could say hydrogen black gas plus oxygen gas goes to water, okay? But if I describe it to you in that way of course you need to translate it from English to chemistry if you will, okay? So the first thing you would have to do is realize that hydrogen gas versus diatomic in its elemental form. So you would have to say, okay, that must be H2, hydrogen gas. The next thing you would do, and so we're going to write the chemical equation for this one, you put the state of matter behind the chemical symbol or the molecular symbol for hydrogen gas. So H2 is hydrogen, G signifies that it's a gas, okay? And then if we looked at oxygen, well we're going to have to remember oxygen was also a diatomic element, okay? And it was a gas, too, like I said, like oxygen gas goes to water, okay? We have to remember water's formula H2O. Of course in this case it would be a gas because it would get so hot. Okay, there are other things about the chemical equation that we are going to have to do to make it correct in describing what we're actually talking about. Of course the next thing would be to balance the chemical equation. We're going to learn about all of this stuff in a little bit. But anyways, the thing you would do is ask do I have the same amount of atoms on both sides? So first I would say do I have the same amount of hydrogen atoms? Well I've got two here and two here, so that's the same. Well what about oxygen atoms? I've got two here but only one here, okay? So I'm going to have to do something to this side. In this case put a coefficient of a two there. And now if I look I've got two oxygen atoms here, two oxygen atoms there. But unfortunately now I've changed my number of hydrogen atoms to this, to four. I've only got two here so I need to put a coefficient of a two there giving me four. So I've balanced that chemical equation. You may have some experience with this if you've already gone a little further in Al but you may not and what we're going to do today and for the next couple of lectures is try to understand all of what I just did in that example. Okay so let's talk about this chemical equation. It's a shorthand notation of the chemical reaction. Shorthand meaning that it only takes a few strokes of the pen to describe exactly what I mean. So it describes all of the substances that react and all of the products that form physical states and experimental conditions. In this case we didn't put any experimental conditions. But let's look at all of these different components to the chemical equation. So the first is going to be the reactants. These are substances that undergo the change in the reaction. We find the reactions on the left side of the reaction. So in this case hydrogen gas and oxygen gas are the reactants. The products are the substances produced by the reaction. So in this case it's going to be the things on the right side. Chemical equations just like we write sentences from left to right. So going from left to right describes your timeline. This is point zero and this is point n. So this is like when you start the thing, this is when you finish. This thing here, this arrow, designates the reaction. So this is what you start with, then it's reacted, then you get your products. So reactants are substances on the left. When two or more reactants are involved in a chemical equation, we have a plus sign between them. This plus sign doesn't mean it's an ion, a water ion. It means hydrogen ion. It just means this reacted with this. So you've got to watch when you have that plus sign that you don't put it in the wrong place and start to mislabel it as what it's not. Again, products are on the right side. If there were two products to this reaction, we would also have separated those by a plus sign. Okay, in this, so the products and the reactants must be specified by chemical symbols. Notice I didn't write the word hydrogen up here. I wrote what it looks like in its elemental form. So they must be specified using the chemical symbols reactions on the left of the arrow, right, as the products. Notice again, here we've got a positive sign or a plus sign separating the products. The physical states are shown in parentheses. In this case, you've got a solid. Mercury 2 oxide is a solid. Mercury is a liquid and oxygen is a gas. This little delta there, that's capital delta as opposed to the little delta that we've been, so it's just a triangle. Notice that it's placed over the reaction arrow. So if we have something placed above or below the reaction arrow like this one is, like that delta it is, those are the conditions of the reaction. So that delta just means that you need to elevate the heat. So if you see that little thing, think of it like a campfire or whatever, so you put a heat into this reaction. And again, notice that it's written above the arrow. So these reaction equations don't always represent reaction. So these chemical equations don't always represent reactions. They can also represent a physical change. So notice here we've got sugar, which is a solid in this little teaspoon here. And then we've got water, which is a liquid in the beaker here. And when we mix those two things up, we form a solution. So they're just intermingled. They're not reacting with each other. So this is kind of a sugar water thing. So what we do is change the state to aqueous, in this case, designating that the things aren't reacting, but they're just found in solution. So not only do these chemical equations represent reactions, but they can represent physical changes as well. So balancing a chemical equation, that's what we did up here with this equation. So a balanced chemical equation is one in which the number of atoms of each element in the reactants is equal to the number of atoms of the same element in the products. Let's just go back and confirm that this reaction is balanced. So we've got two atoms of hydrogen, but we've got two times those two atoms. So there's going to be four hydrogens there. We look over here. Water has two atoms of hydrogen in it, and we've got two water molecules. So there's going to be four hydrogens there. Go back to over here. Two oxygens, because each oxygen molecule's got two oxygens in it. And here we've got two oxygens as well, because each water molecule has one, and we've got two water molecules. So we've got to balance the equation. The reason we have to balance the equation, or there's a number of reasons it makes our lives a lot easier when we start doing calculations using the chemical equation. But the physical reason why we have to do this is so we can obey the law of conservation of matter. If you recall what we were talking about when we talked about the law of conservation of matter, is that matter can't be created or destroyed in any chemical or physical process. And if that's the case, then we have to have the exact same number of atoms on the left and on the right. So that's why we balance the equation. And again, it'll make your lives much easier when you say, well, if I have .656 moles of hydrogen in gas, how many moles of water do I need? If you don't have it balanced, you can't do that problem. In this case it's a gas, or in this case it's a gas, this case it's a liquid. I don't think you need to bog yourself down about that. It's just depending on what temperature you do the reaction at. Above 100 degrees Celsius, water's a gas, but below it's a liquid. So they must have done this reaction up there, at below 100 degrees Celsius. And I did my reaction here at above 100 degrees Celsius. So you've got to remember that molecules can change physical states. So it all depends on what, how much energy you're pumping into the system. So don't bog yourself down about that. It'll tell you what it is on each chemical equation. So these things that we put in front, these numbers that we put in front of these molecules here, they're known as coefficients. And what they represent here. So let's just go back into our chemical equation. So what does this represent? This represents what we've got. One, two hydrogen molecules that react with one oxygen molecule to form two water molecules. So that's exactly what this is saying. It's the same thing as saying, of course, you know, like if I have two of these they react with one and form two. But if I have two moles of this, it reacts with one mole and forms two moles. Because remember, mole is just a number like a dozen. So if I had two dozen of these, that would react with one dozen of these forming two dozen of these. So if I had 0.5 of these, 0.5 moles of this, then this would be 0.25, this would be 0.5. So you can do that through the chemical analysis or the molar and moles and multiply that by the molecular weight or divided by the molecular weight. We'll give you the number of grams you're going to do all of those calculations with balanced chemical equations. So again, we talked about the law of conservation of matter. Matter can't be gained or lost in the process of a chemical reaction. Therefore the total mass of the left side and the right side have to be equivalent coefficients of what you use to do the heat and energy are the same thing. So like we were saying, the chemical equation shows also the molar quantity of reactants. It doesn't only show the individual particles that are reacting as just the molar quantity, just like what we just said. So it's the relative number of moles of each product in reactants as indicated by placing this coefficient in front of it. So I have a relative number of two moles of mercury-2 oxide that form relative two mercuries to one oxygen. If there's no coefficient in front of the molecule in the chemical equation, then it's understood to be one. So here's a chemical equation that is unbalanced. Let's go ahead and try to balance it. Again, that'll always be given, so it's not something that you'll have. So the way I would do any reaction that has carbon in it is to try to balance our carbon first, then our hydrogens, and then our oxygens. So this is the way I would always do it. So the first thing I would ask myself is on this side of the reaction, how many carbon atoms do I have? Three. And how many do I have over here? Just one. So I'm going to have to put a what? A three in front of carbon dioxide. So now I have three on that side and three on that side. So that's cool. Remember, we don't want to put a one there because that's implausible. So now let's go to hydrogen. Well, how many do I have here? Eight. How many do I have here? Two. So I'm going to put what here? Four. Okay, so we've now balanced our carbon and our hydrogen. The way I would do oxygen is to come over here first, and then balance it here, and then go over there. Or figure out how many you've got here, then go over that side and balance it. So we've got how many oxygens altogether do we have? So we've got four plus six, right? It's ten. And how many do we have here? Two. So we're going to have to put a five there. So make sure you add up both the six that's here and the four that's there. So that one wasn't so bad. Let's try a different one. That might be a little bit hard. So instead of C3H8, let's do C4H10. For all of those of you who are confused about this, we're about to do one more. So let's do it together. Okay, remember you want to balance C, then H, then O. Remember this is different than what was up back there. Balance C, then H, then O. If you've got just those atoms in the reaction equation, and in this case you do. So the first thing you're going to balance is C. You've got four here and only one here, right? So you're going to put a four there. Everybody got that far. What about hydrogens? You've got ten here. You've got two there. So what are you going to put here? Five. And then what did I tell you? You've got to go to the right side of the equation, figure out how many oxygens you've got. Okay, so how many do you have? Five times one is five. Four times two is eight. The whole equals 13. Okay, so how many do I need to put here? What do I do now? Decimal, can you have half of an oxygen molecule? No. So this is why I wanted to do this problem. So what you're going to do is actually, you could start with the decimal. So 13 divided by five, right? That's going to give you, sorry, 13 divided by two or 7.5 would give you the right amount, right? So what you're going to have to do, since you can't have that 13 divided by two, you're going to have to multiply all of it by two. Okay, so if I've got 13 divided by two here, I've got one here going to multiply the whole thing by two. When I do that, two times one is two, right? Two times 13 over two is 13. 13, right? Two divided by two is one. So two times four is eight, and two times five. And now is that a balanced equation? It should be, right? Okay, because we had a fraction here, right? 13 divided by two. Okay, we can't have half of a molecule, right? It's like saying, okay, this is in our reaction, but we want to cut it in half. That doesn't make any sense, okay? So you can't do that. So it was, the coefficients were one, 13 over two, four, and five, right? So since we can't have this one, we got to multiply the whole thing by two. Okay? So watch out with those types of reactions. I can only imagine something like that would be on the table. There's that first reaction that we balanced doing it in the same fashion that we did this one. What I'd like you guys to do on your own is balance these equations, and then if you have any problems with any of them, let me know, come to office hours, let me know, and we'll balance these together. Okay, so we discussed this already, but this table really does, this table isn't in your book unfortunately, but it really does emphasize what kind of information the chemical equation is actually telling us, okay? So it's telling us a lot of information, even though, you know, it's not always so obvious or right apparent when looking at the chemical equation. Of course, it's telling us in this equation here, it's saying two molecules of this, this butane stuff, so whoever lit a cigarette lighter, this is the chemical equation that you've used, okay? Two of these molecules react with 13 of these molecules to form 8 of these molecules and 10 of these molecules. That's what that says, okay? That's probably the most obvious thing that it's telling us, okay? And that's kind of described pictorially up here, although for a different reaction. But also, we can think about moles, because these are the relative amounts of things. So for every two moles of this, we need 13 moles of this to react, and 8 moles of this will be formed, and 10 moles of this will be formed, okay? So that's also what it's telling us. And that might seem obvious to you too. But it's also telling us that the total weight over here is equal to the total weight over here. So for this equation up here, it says while the total mass is 204.09 grams on this side, will it better be 204.09 grams on this side, the molar mass that it is? If we're looking at AMUs, it's got to be 204.09 AMUs on this side, 204.09 AMUs on this side. And notice, you can figure out the masses, remember converting AMU to mass, and a lot of you had trouble with that on the test. I would suggest that you make sure you guys know how to do that, because it's going to be just coming back, and coming back, and coming back. So if you weren't able to convert from moles to grams, AMUs to moles, that's all going to be the rest of the class, okay? So you just got to figure it out. You got to figure it out. But the chemical equation can tell you all of that information. And with those three simple figures, you can get even more and more and more information about these sort of reactions. Okay, so we're going to talk about a series of reactions. The first reaction we're going to talk about is a type of reaction. It's called a redox reaction. And it's the combination of two words, this word redox. It means it's squished together these two words, reduction and oxidation. And we're going to learn about oxidation numbers, and they really provide kind of a convenient way to keep track of which molecules are reacting in what way. Okay? So let's go ahead and erase. Can I erase this equation? Everybody's got it? The way that I learned this, way back when, maybe even in high school, was through this two acronyms. There was a phrase also. It was Leo de Lyon says ger. But Leo describes oxidation because if you lose electrons, you get oxidized. So that's Leo de Lyon. He's saying ger. So there's going to be, so there's all of these other ways to look at it. I really think this is the easiest way to keep it straight. Another couple set of acronyms is oil rig. But it's kind of more appropriate for the area that we're in, but it doesn't, I don't know, oil is, oxidation is losing. So it doesn't really describe the electron. So I think this Leo de Lyon says ger might be the best way to think about it. Okay? So if we look at this reaction up here, this is an oxidation reduction reaction. Any oxidation reduction reaction deals with the transfer of some amount of electrons from one species to another. Okay? In this case, zinc as its elemental form, anything in its elemental form, let's look at a few things that may be in their elemental form, zinc solid form. Its oxidation number will be zero. Anything in its elemental form will be zero. Okay? Like chlorine gas, right? That's going to be Cl2 gas. That's also zero oxidation. The oxidation number can be either zero or a negative number or a positive number. Okay? If it's a negative number or a positive number, it will equal the charge of that ion. Okay? So for group one elements, what's the common charge that a group one element will attain when doing a reaction? Positive one, positive one. So that's the oxidation state for all group one elements. Okay? So oxidation state and ionic charge are kind of, what do you say, interchangeable terms. Okay? So if we look at group two elements, what's the most common number for those? Plus two. That's going to be their oxidation state. Okay? Normally, these group sevens will be negative one. So that will normally be their oxidation state, although every once in a while you'll see them with different oxidation states. And then these guys will normally be negative two, these will normally be negative three, but occasionally the non-metals have very strange oxidation states. But we'll figure it all out. Don't worry. So if we look at these two reactants up here, what's the oxidation state of copper here? And this, and this reaction here. Two plus, how did you figure that out? Because it looked at it and it said two plus next to it, right? So what's the oxidation state of zinc here? Zero. Zero, how'd you figure that out? Yeah, because it's in its elemental state is what you want to think about, okay? So what happened here? How many electrons is this missing here? Two electrons. What happened from here to here? What did it do? It must have gained those two electrons, right? Because it went from oxidation state what here? Two plus to oxidation state what here? Zero. Zero. So how do you get from two plus to zero? You have to get two negative charges, right? For every one negative charge, you can assume that's one electron, okay? So it's gained two electrons. What happened to zinc here? So it went from a what oxidation state? Zero to over here? Two plus. So what did it do? It lost electrons, right? Because if you lose negative charge, you gain positive. Okay, so which one of these things got oxidized and which one got reduced? Which one got oxidized? Oxidized. So which lost electrons? Did copper lose electrons? No. No. Uh-uh, zinc lost electrons. How do you know that when you look at the oxidation state? It'll tell you, right? It in fact lost two electrons, right? What about copper? What did it get? Reduced or oxidized? Copper did what? Now, gained electrons, okay? So remember, Leo's alliance is Ger, right? Okay. So if it gained electrons, what happened? It got reduced, right? Okay. So you gotta remember this in relation to oxidation numbers. Okay, that was a crash force in oxidation numbers. We're gonna go over it quite extensively. So what's happened here? This case is this zinc atom transferred two electrons to the copper 2 plus here, forming these two products. Okay? This is the redox reaction because you have that transfer of electrons. So here are the rules you want to remember for oxidation numbers. Any uncombined elements' oxidation state is zero. Just like what we said here. The zinc in its elemental form is zero. Fluorine in its elemental form is zero. If I had H2, is that the common elemental form for hydrogen? H2? Everybody should be like screaming at me, yes, right? If you don't know that, that's something you need to know. Okay? You need to know that. So this elemental form, so it's going to be oxidation state zero. So most of the elements are just going to be kind of like zinc here. Where it's not a diatomic element. Okay? So all the metals will be that way. So it's just the non-metals you've got to worry about. So a simple lyon, like what we were saying, group 1's are positive 1, group 2's are positive 2, group 7A's are usually negative 1, usually negative 2, and so on. Okay? For example, magnesium 2 plus is going to be 2 plus observations. Oxygen 2 minus is going to be a 2 minus oxidation state, and chlorine minus is going to be a minus 1 oxidation state. So when they're in compounds, group 1 and group 2 will always be plus 1 and plus 2. Hydrogen in a compound will always be plus 1. Okay? So this starts to give you weird oxidation numbers. We'll do this H2SO4 and figure out what the oxidation number for sulfur is in that. Okay? But another rule I want you to remember is the oxidation number for oxygen is always negative 2. Okay? So oxygen's the one that doesn't change. Except in peroxides. We won't see any peroxides in this glass. You might see them in some of the allocinins, but not in any of the tests. Okay? So the last thing you want to remember is the oxidation numbers of all the atoms in a compound equals 0. Okay? So let's go back to H2SO4 and figure out the oxidation numbers now that we know all these rules. So if I give you a compound like H2SO4, you should be able to figure out what the overall oxidation numbers of each of these atoms are. So the first thing you want to remember is that the overall oxidation combined should be 0. Okay? You also want to remember in compounds, hydrogen is always plus 1. Okay? So how many hydrogens do we have in this compound? Two. So it's going to be 2 times plus 1. Okay? We're going to add that to something, whatever sulfur is, because we don't know yet. Okay? But oxygen is always going to be what? 2 minus. 2 minus. Okay? So those are two of the rules you remember, hopefully. So how many oxygens do we have? Four. Four. Right? So it's going to be 4 times 2 minus. Okay? So what is this number altogether? Negative 8, right? Everybody agrees with me, hopefully? What's this number altogether? Two. So 2 plus 1 minus 8 equals 0. That must be 6, right? 6, right? Is that right? So for those of you who are confused, this is just addition, right? We're not doing anything real special. Okay? So it's going to be 6 because 2 plus 6 equals 8. 8 minus 8 equals 0. Okay? So normally you would expect the sulfur atom to have a negative 2 oxidation state, right? But in this compound, sulfur's oxidation state is actually plus 6. Okay? You've got to watch out for those. The nonmetals will do this. Oxygen won't do this because oxygen is always minus 2. Hydrogen won't do this, it's always plus 1. But all the other nonmetals, you've got to watch out for. Okay? So you always want to do your hydrogens and your oxygens first whenever you're trying to figure out the oxidation number of a compound. Okay? Let's stop there for today and hand back the test. So please don't leave before you get your test back. And we will pick up with figuring out the individual oxidation numbers of atoms and polyatomic ions.