 Okay, guys. If you don't mind, we can get started. Today we'll have a short lecture. There's a sign-in sheet going around. Make sure you guys sign in. I might do, I haven't I might do another kind of impromptu review session going over that quiz for people. I have a video of it, and I haven't posted it yet because there was a misprint on the key and I wrote something incorrectly. So I'm hesitant to post it, you know, because I don't want you guys to think that one of these answers was correct. But I might do another one for anybody who wants to show up. And we'll talk about that. Maybe I'll send you guys an email or something tonight about it, okay? Anyways, so I think the last thing we were talking about was equilibrium. So hopefully we can finish up chapter 8 today. Rules for writing the equilibrium constant. Remember the equilibrium constant is big K. So remember that equilibrium constant expressions can only be written after you correctly balance a chemical equation. Okay, so if you have a chemical equation that you're trying to write the rate constant for and if it's incorrectly balanced, you won't be able to write the rate constant. You'll get the wrong rate constant. Each chemical reaction has a unique equilibrium constant of unique K at each specified temperature. So if you change the temperature of the reaction, K is going to change, okay? Because the concentrations of the reactants and products will be different. You don't really have to think too much about that. Any problem I'll give you is going to stay at the same temperature. It's just something good for you guys to remember. The one thing you do or one thing, one of the many, many things in chemistry that you want to remember, is that brackets around a substance represent molar concentration. So remember moles per liter, right? So that's like chemists' favorite concentration unit. That's why we've hit on it so many times, okay? And so whenever you see brackets around solution or something like that, that means it's in molar concentration. Equilibrium constants themselves are unitless. Remember, you're going to have concentrations over concentrations, and sometimes with your canceling of units, you may get some units in your answer. But you want to remember, even if you do get units in your answer, you just want to drop them, okay? Because K is a unitless number. And the only things that go into K, the evaluation of K, are the concentrations of gases and substances that are in solutions, okay? So pure liquids and pure solids aren't shown in the equilibrium constant. So let's just write the equilibrium constant for these other two reaction equations. So notice the equilibrium constant for the first one is the molar concentration of ammonia squared. Divided by the molar concentration of nitrogen times the molar concentration of hydrogen cubed, right? So where did we get these numbers from? That molar concentration of ammonia comes from ammonia being a product. And it's in the equilibrium constant because it's a gaseous product, okay? Notice that we have a little two superscript up here, a squared, right? That comes directly from the coefficient here of ammonia. So recall, the products go on top here. So since ammonia is the only product, that's going to be the only thing that's on top. And then the same rules apply for the bottom, but the bottom being the reactants. And you multiply them together so we got concentration of nitrogen times the concentration of hydrogen cubed, right? Three superscript is cubed. So let's try the next one. So first things first, we want to make sure this reaction equation is balanced. Is that reaction equation balanced? As far as I can tell, yes it is. Okay, so let's write the equilibrium constant. So K eq equals, it's always going to be something divided by something else, okay? So it's going to be concentration of Hf, right? Squared. Everybody got that, right? And the other two would just be concentration of H2 times the concentration of F2. Remember, brackets mean molar concentration, okay? So let's try the last one. The last one incorporates different types of, different states of matter, excuse me. So remember, the K eq is always going to be the products over the reactants. So the products, we're going to say manganese 2 plus, right? Why would we can do that one? Because it's in solution, right? Everybody sees that it's in solution. And then chlorine gas, right? So we can multiply that times chlorine gas there. No coefficient. And water, will we put that into equilibrium constant? No, because it's a pure liquid, right? Okay? So manganese oxide is a pure solid. So we're not going to put that into equilibrium constant. Those protons are hydrogen ions there. Those aren't pure solid or liquid, so they go in, right? But since they have a coefficient of 4 there, we're going to have to make this concentration of H to the fourth there. And then chlorine ion, same thing, CO minus squared. So hopefully everybody got that as the equilibrium expression. Okay, remember, a reversible reaction is allowed to proceed until the system reaches equilibrium. Sorry, this is showing an equilibrium reaction proceeding until equilibrium. So you can see the concentration of the products or the reactants here are going down, down, down, down, down. And the concentration of the products are increasing, increasing, increasing. Does everybody see that from this graph? Notice that once you get past this point here, some time point, if I can keep the thing on that exact point, the concentration doesn't change after that point. So after that point, the amount of the reactants and products no longer changes. So the calculation of K is actually the analysis of this reaction mixture to determine the molar concentrations. Okay, so now we know how to make the expression for K. Well, if I gave you equilibrium concentrations, you should be able to figure out what K is. Okay, so let's try this problem together. So first off, let's write the expression for K. So KEQ equals concentration I squared times concentration H squared, hydrogen squared. What am I saying? Concentration iodine times concentration hydrogen divided by hydrogen iodide squared. Notice I give you here the concentrations and molar units. So you can just stick those concentrations into this expression. So H2 is going to be 1.72 molar times I2 1.72 molar divided by HI 0.54 molar. So really, I don't need to put these units in here, okay, because the equilibrium constant is going to be unitless. Yeah, what's up? Oh, I just talked, yeah, I just said that, yeah. Yeah, it does not matter. This doesn't matter, because if I multiply 3 times 4, right, it's the same as if I multiply 4 times 3. So don't get hung up on that stuff. Okay, so let's try to figure this out. So all we really have to do is plug these numbers into our calculator, 0.54 squared and then 1.72, well, squared, right? Okay, I got this number 10. Well, that's it, right? So I can only get up to two significant figures there, right? Why? Because my one concentration has two significant figures. So in other words, my KEQ equals 10, okay? So let's, don't let that number without a unit, I know it looks weird, right? But don't let it throw you off, okay? Because K has no units. We can understand a lot about this reaction from its equilibrium constant, honestly. The one thing I want you to definitely be able to see, like right away, is that of course the equilibrium constant is greater or less than 1. It's greater than 1, right? Is it greater or less than 1? Greater than 1. Greater than 1, okay? So do you guys remember what we talked about on Friday, if it's greater than or less than 1, what it means? What does it mean if the equilibrium constant is 1? Does anybody remember? Yeah, so the concentration of the reactants is equal to the concentration of the products, if it's 1. Okay? If you guys recall if it's greater than 1, what it means? It's mostly products, right? So what does that mean? What does that inherently mean for the reaction? It's mostly products. Does this reaction... Not really so much like that, but this is a good reaction, right? This reaction will go forward. Okay? Right? If my equilibrium constant was 0.0000001, is that greater than 1 or less than 1? Less than 1, right? Less than 1. So does that mean there's going to be a lot of reactants or a lot of products? A lot of reactants, right? So if there's a lot of reactants, is this reaction very useful? Am I getting out of this reaction what I want? No. Why not? Why not? What do you want out of a reaction? More than likely if you're doing the reaction. Products, right? So if your equilibrium constant is lower than 1, it's probably what we call a bad reaction, right? Because it doesn't go, it'll give you what you want, right? Your equilibrium constant is greater than 1. It's probably a good reaction, right? Because it's giving you mostly products. In fact, this reaction is giving you 10 times the amount of products that there are in solution as reactants. That's what that equilibrium constant is telling you, okay? So you want to think of it in terms of good and bad reactions. You always want to equate it to 1, okay? Equilibrium constant of 1 has the same amount of reactants and products on both sides. So now let's talk about Le Chatelier's principle. So this guy Le Chatelier, some French dude who's been dead for a long time, so we don't have to really worry about it, right? But this principle is very, very important, okay? It really tells you which way a reaction's going to go, okay? And you can predict it. So when a chemical system is at equilibrium, we can disturb that equilibrium, right? You can imagine, like, what if you got your chemical system to equilibrium and you started taking out some of the reactants or products, right? Concentrations would change. The relative concentrations would change. And your K wouldn't be the same anymore, right? And your K always is the same at the same temperature in the same reaction, okay? So you got to watch that. So if you remove some products or remove some reactants or add some products or add some reactants after you think that equilibrium, it's going to adjust again to get back to equilibrium with that same K, okay? Does that make sense to anybody? Hopefully. Hopefully, after our discussion today, if you go home and read this stuff, it'll kind of click, okay? But anyway, so, yeah, Le Chatelier's principle says if I've got this situation where my, you know, concentration of reactants, you know, decreased, decreased, decreased over time and then stayed there, right? And then my concentration of products increased, increased, increased over time and stayed there, right? What would happen to this system if I put like a ton more products into it, okay? It would, of course, make the product concentration increase very high, right? But our K is going to be the same, okay? Because we haven't changed the temperature of the reaction. So K is going, it's going to be 10 still if it's this reaction, right? It's still going to be 10. So if I increase this product, right? That means that I increase this very dramatically. This part very dramatically. That means K would increase very dramatically, too. In fact, it wouldn't be called K anymore. We call it Q if it's before equilibrium is reached, okay? So like this whole part, we call K Q, Q, Q, Q, Q until it gets to there and we call it K. But you don't really have to worry about that too much, okay? You just want to know that if you make this very big and this has to be 10, then you're going to have to adjust this number to make more products or make more reactants, okay? So this is what happens when you add one or the other, okay? This is Le Chatelier's principle. So the things that will disturb the equilibrium are a change in temperature. Like we said, that'll change K. A change in pressure, ironically enough, will do this. That's because we're dealing with gases in equilibrium constants. And a change in concentration. Any one of these three events will change your reaction direction, okay? So let's look at a generalized reaction. Or actually, we can just keep looking at this reaction where we've already solved K for. That one over there. So let's write up the reaction really quick. What was it? Okay, so that's our reaction there, right? So if this is our reaction and reactants are... Well, so it's an equilibrium reaction. It's actually going to be like that, right? Because it can go back and forth. So if it's at equilibrium and reactants are added, so if we add that, then the reaction's going to be pushed forward like that, okay? If we add this product, it's going to be pushed the other way, okay? So we can push the reaction back and forth. So if we add this product, we'll also be pushed back toward the reactants. So whatever you're adding, it's going to go the opposite way. So that's what this is stating here. If reactants are added, the position of equilibrium will shift towards the products. And that's what we showed first. If the reactants are added, the equilibrium position will shift towards the products, okay? In other words, we can think of this arrow going back and forth at equilibrium. And then if we add this, we have big arrow there, little arrow there. If we add reactants, okay? If we add products, big arrow going this way, little arrow going that way, okay? If reactants are removed, right? So what happens if we remove reactants? So if instead of putting more in, we took it out. What do you think's going to happen? Well, we're going... things are going to realize that there's not enough reactants to make the K what it's supposed to be. So they're going to convert from products to reactants. Okay? So if I remove something, the other thing will convert to it. Because remember, if I put these two guys together, they make this thing, okay? So they can transform back and forth, back and forth. If I remove this thing, what do you think will happen? The equilibrium will shift, not towards the reactants, right? It'll shift towards the products. Okay? So you... some people have a hard time with this at first, but really it's just the same thing over and over and over with every reaction, okay? So this is just one of the things. This is if we're adding or removing concentration of products or reactants. Remember, if we're adding here, it's going to want to get rid of the excess, so it's going to convert some of them to the products, okay? If we're taking away here, right, these things are going to be like, oh crap, we need to have more reactants in here, so we're going to make more reactants and go this way. If I remove this guy, this thing's going to be like, oh no, you know, let's make some more. Or if I remove this one, same thing, okay? So that's Le Chatelier's principle in relation to concentration. Here is exactly what we were talking about with a different reaction, okay? You can see the equilibrium constant at 523 degrees Kelvin for this reaction is 24. What does that mean if the equilibrium constant is 24? Is that a good or a bad reaction? How about that? That's a good reaction. Why do you say that? Because it's bigger than one, right? It's bigger than one, okay? So this is exactly what we were talking about. If you increase this, it goes to the right. You can see the green, right? If I decrease that, it also goes to the right. Okay, so study this little diagram. Whenever this is Q, like I said, it's not something that we're really learning in this course, but don't get confused by it. Before the reaction reaches equilibrium, okay? So you can see here, again, effective changing concentration for removing it. It's going to go backwards. You could calculate the equilibrium position if you wanted by doing this, essentially, doing this sort of analysis of it. Let me come back to this. I don't know if we'll go over... I can't remember if we did this last spring. If we'll do it for the test, then I'll come back to it. Okay, let's skip over it for right now. I just want you to know qualitatively what goes on if you add or remove stuff, okay? For right now, I don't know. Does any of that Al problems deal with anything like this? Has anybody gotten there yet? Okay, so not that you know of. If they start dealing with it, then we'll talk about it. How about that, okay? But right now qualitatively, I don't think we need to go this in depth into it. Okay, but this, we should be able to do. So here's our reaction, right? Sulfur, hydrogen sulfide, sorry, oxygen gas go to sulfur and water, right? Well, what happens to the concentration of water if oxygen is added? So if we add this stuff here, what happens to the concentration of this? Yes, can you figure that out? It's going to get bigger, right? Because we're adding this, so the reaction's going to go to the right. Does that make sense? Okay, and what about if... what about the concentration of oxygen if we remove hydrogen sulfide? If we take some of this away, what's going to happen to the concentration of oxygen? Is it going to get bigger or smaller? So if we remove this, what's going to happen to these arrows here? If I remove a reactant, what happens to the arrows? Just because everybody said some answer, don't make it make yourself believe that that's the right answer, okay? So what happens if I remove a reactant? What is the arrow? It goes to the left, right? So the concentration of this is going to do what? Increase, right? Increase, right? Because if this is going to the left, right, this is getting bigger. What about the concentration of hydrogen sulfide if oxygen's added? So if we're adding here, what is the reaction arrows doing? We're going more to the right or to the left. So the right. So is this decreasing or increasing if we're going to the right? Decrease it, okay? So I know that half of you are like, oh, this is so easy, and half of you are like, crap, I'm getting it backwards every time, right? So if you're one of the half that's getting it backwards every time, you just need to change your thinking around because everybody was like, yeah, this is easy, easy, easy. So you got to remember that it is easy. You just got it backwards, okay? So go back and study these diagrams that we've been talking about, okay? And then there's quite a detailed analysis of what's going on here with relation to the equilibrium constant. So you can see it going up and down. Here's another one I'll let you do on your own. And the last thing we'll talk about is the change in pressure. Remember we said change in pressure can increase or decrease concentrations through Le Chatelier's principle. But this is only if we've got gaseous particles in our reaction, okay? So if all of them are in solution, like if we had H plus aqueous plus OH minus aqueous, we'll say whatever. Well, this is not a very good one, but we'll do it this way. Aqueous, aqueous, okay? So if we have all of our stuff being aqueous, changing the pressure doesn't affect anything because changing the pressure only affects the gases. So the way we can change the pressure is we can change the pressure three different ways. We can change the concentration of the gas as we've been doing already. You can add a gas that doesn't react with the other reactants or other products called an inert gas. So you can imagine adding a noble gas like neon, something like that that won't react with anything. Something that won't react with other things. Yeah, so like a noble gas would be a good example of an inert gas. So like neon, helium are gone. So again, adding an inert gas has no effect to the equilibrium position because you're not changing the concentrations. Or you could change the volume of the reaction vessel, right? You could imagine taking the same number of moles, putting them in a bigger container, and that changes the pressure, right? It's a lot less pressure. It's like going into outer space or something. Field trip. You know, I read, I heard about the first astronauts, they did a bunch of experiments on them, kind of talking about Le Chatelier's principle, that they fed them a bunch of beans, right? And the ones that would make the least gas after eating the beans, right? Those are the ones that got to go up because like if you decrease the pressure, right? The gas would expand, right? And it would explode your insides, right? So they wanted to make sure that they wouldn't... Well, I mean other things produce gaseous products, you know? I'm just saying, yeah. So they figured out which ones could produce the least of these things because of Le Chatelier's principle because when you go out to outer space, right? The pressure decreases quite rapidly and things expand quite dramatically, right? So you can imagine you got a bunch of gas in you, you decrease your pressure, and then you become like really big, you know? But that's what happens, right? So chemistry is everything, right? You got to watch out, you know? Even if you're going into outer space, right? Okay. So the things we need to remember or know is that changing the volume of the reaction vessel causes a major change in the equilibrium position. So when the pressure increases, the volume decreases, or the volume decreases, the system decreases the moles in the system, okay? So this is like if we have this reaction here. This reaction here actually, gas, gas and gas, right? So on the right side of the reaction, how many moles of gas do we have? If this is a correctly balanced equation. Two moles of gas. Well, this is not a very good one because there's two moles and two moles. Let's do one that has three. So if this has 3A gas, it goes back and forth to 2B gas, okay? So on this side of the equation, we've got two moles, right? On this side of the equation, we've got three moles. Three moles takes up more space than two moles, okay? So if I decrease the pressure, it's going to push the reaction that way, okay? All right, sorry if I increase the pressure. Sorry if I decrease the volume. But if I decrease the pressure, like with the astronauts, right? It's going to push the reaction this way and make them blow up, okay? So it's decreasing the volume and pressure, okay? Or pressure and volume, so you've got to watch out. So when the pressure increases or volume decreases, the system wants to decrease the moles of gas, okay? If the pressure decreases or the volume increases, the system increases the moles of gas to occupy the remaining space. When the volume doesn't change in the course of the reaction, changing the volume doesn't do anything, right? So you can see here on this equation, which side? So this side has how many moles of gas? None, right? This side has how many moles? One. So if I decrease the pressure of this reaction, which way is it going to go? So we say, so the right is forward, so the left is reverse, okay? So if I decrease the pressure, which way is this reaction going to go? Decrease the pressure. If I decrease the pressure, which has more moles of gas? This side or this side? The right. So if I decrease the pressure, which way is this going to go? To the right. So the right, right? Decreasing the pressure means what? In relation to volume. The increase of volume, just like we talked about with the astronauts, right? So the Earth relative to space, which one's bigger? Which one's bigger? Earth or space? Space is bigger. So what did I say about the astronaut if he eats too many beans and goes into space? What happens? Does he get smaller or get bigger? Okay, so now you can remember, okay? Okay, the pressure is decreased in outer space or increased? Decreased, okay? So now you remember all of Le Chatelier's principle. Remember the exploding astronaut, okay? Try to do the rest on your own, okay? There's notes there. I would go over more, but you guys got to do evaluations. So when you guys are ready to do evaluations, there's one page here. You've got to do one page here. I've got to assign somebody who wants to... Do you want to do it? Okay, Aniz, there's this instructions. Instructions that you have to read if you don't mind. Guys, if you could for one second. So if you don't mind on these evaluations, remember this evaluation is only for 1406 lecture, okay? So don't talk about your lab. Don't talk about good or bad things about that. You'll have time to talk about that stuff, okay? Also, if you don't just be like, Heath is so awesome or Heath really blows, you know? Because that doesn't give me any feedback, okay? So if there's something that you like or don't like, kind of expand on it. That would be the most appropriate way to do this, okay? Also, I brought... My wife works at Devereux Gardens, and she gets a lot of brownies a lot of the time, and we had to take... We didn't have enough room in our freezer to keep some of them, so... Since it was evaluation day, I figure. I'll bring you guys some brownies. But there's some napkins there. Feel free. I cut them up into little pieces, so if it looks kind of... They looked mangled. It was because I tried to cut them up evenly, but it wasn't able to. But please feel free to take...