 Okay guys, if you don't mind, get started with class. We'll have a short lecture this morning. And then get on with the quiz. There's the sign-in sheet. So of course, this didn't start at the right slide. So I believe this is where we, guys and gals, I know we're all studying right now. But we'll just do a little bit of a lecture. I think we finished here last time. So if we look, we can look at these two different graphs. And we can see two different things about equilibrium. So initially, the forward reaction is rapid. So the reaction, this is the concentration or the rate of the reaction, the forward rate of the reaction, goes quickly, quickly, quickly when the reaction concentrations are high. The product concentrations are negligible. You can see here they have zero concentration. And they increase as the reaction progresses. So the rate increases as the reaction progresses. If you notice here, the rates will become equal at the time the reaction gets to equilibrium. So the rate of the forward reaction is equal to the rate of the backward reaction. So you can also see the concentrations of the product increasing quite significantly until equilibrium is established. The reactants decreasing quite substantially until equilibrium is established. And then the concentration stays the same of both of them. Remember, even though the concentration stays the same, right, the rate still changes, or the rate still, the reactants still go back and forth from reactants to products. So the rate of the forward reaction is equal to the rate of the reverse reaction here. So the reaction continues to be dynamic, but no change is observed due to the rates being equal. Okay, so let's introduce this concept of equilibrium constant. So equilibrium constant is big K. Okay, so this is a number that's equal to the particular ratio of equilibrium concentrations of product and reactant at a particular temperature. So what you find is that the magnitude of K indicates how far the reaction will proceed towards product. So if you can look at these three boxes, they represent three reactions that have gone to equilibrium. Notice here, in the first reaction, in reaction A, the reactant is blue, the product is red. At equilibrium, we have a bunch of reactant and very little product. So what you find is that you got a small K when this situation arises. Small being less than one. When the opposite situation is true, at equilibrium we've got a bunch of product but very little reactant. That means we have a large K. This should say mostly reactant and mostly product, I guess. And then when we have intermediate value, that would be K equals one. So this means that we have exactly the same concentration of products and of reactants when your K equals one. So here are some ways to analyze or some reactions that have been analyzed using K and using these reaction concentrations and rate diagrams. So you can see here the equilibrium constant, N2O4, goes to 2N02. So what we say is the constant of the forward times the concentration at equilibrium equals the constant of the reverse reaction times the concentration at equilibrium. Notice we square that unit there. Why do we do that? Because we've got a coefficient of two here. What you find is K is the... So let's just take from the one equation and make the other equation. So what do we got? K forward is N2O4. That equals N2O4 at equilibrium. This equals K reverse times the concentration of N02 squared at equilibrium. So that's what's written there. What big K actually is, that's going to be the constant of the reverse reaction over the constant of the forward reaction. So we'll set this equal to that. Sorry, do I have that wrong? Yeah. So when we set this equal to that, we're going to get K equals that. So let's put these things on the same sign. So we'll divide both sides by the K of the reverse like that. So now we have K forward over K reverse. And remember that equals big K. And we also still have this times N2O4 at equilibrium equals N02 squared at equilibrium. And then we can divide both sides by N2O4 at equilibrium. So we get this. Big K equals K forward over K reverse equals the concentration of N02 squared at equilibrium over the concentration of N2O4 at equilibrium. So you don't really need to know how to manipulate this equation. Just know that you're going to put the products with their coefficient as a superscript over the reactants with their coefficient as a superscript. Notice in this case, we don't have a coefficient for N2O4, so there's no superscript here. So K is going to equal overall. We've got it written over there, but just to make it more obvious is the concentration of N02 squared over the concentration of N2O4. So this is how you're going to get K for everything. And again, you're not going to have to go through this manipulation every time. You can just do K equals this stuff here. So remember, equilibrium constants aren't just for reactions. They can be for phase transfers too, just like here, sugar solid going to sugar aqueous. Well, the constant, the equilibrium constant for that reaction is the concentration of the aqueous sugar over the concentration of the solid sugar. So this is the general equation here for the equilibrium constant. When we got A, A plus B, B goes back and forth to C, C plus D, D. We see that KEQ equals the concentration at equilibrium of C to the subscript C. Remember, it's the products over the reactants here. Concentration of D, superscript D. Concentration over concentration A, superscript A, that's coming from the coefficient. Concentration B, superscript B coming from the coefficient. So the exponents in the rate expression are numerically equal to the coefficients. And then, so there's some rules that you need to follow. Remember, the equilibrium constant expressions can only be written after a correct balanced chemical equation. If you don't balance it correctly, you won't be able to get the superscripts, okay? This, yeah, each chemical reaction has a unique equilibrium constant value at a specified temperature. So if I tell you the equilibrium constant is this, at zero degrees Celsius, if I change it to 20 degrees Celsius, it's going to be a different equilibrium constant, okay? We don't have to worry about that too much. I'll tell you the equilibrium constant that we're using. You're not going to have to manipulate that. Remember, the brackets represent molar concentration, okay? So you're going to have to get all your concentrations and molar values, okay? If they're in, like, grams per mill... grams of stuff per milliliter of solvent or, you know, something else, you're going to have to convert it to moles per liter, okay? So if they're in, like, weight weight weight or weight volume or something. All equilibrium constants are unit lists, okay? So this is where I know throughout the class I've been telling you that you can cancel out all your units, okay? But when we do equilibrium expressions, what you'll find is that the units won't cancel, okay? So sometimes they will, sometimes they won't, right? If you can imagine this one right here, we said what? K is unit list, right? But if we look at the units here, right? What's the units going to be in the brackets here? What do we just say? Brackets tell us what? Molar, right? Good job, guys, okay? So molar squared, right, is what our unit should be. Does everybody see that, right? What is the units down here supposed to be? Molar. Is it molar squared? No. So what are we going to get out of this? We should get molar, right? But we don't get anything because it's unit list, okay? So you've got to watch out for this, okay? So this is the first example that you've seen, again, right? The first example that you've seen that you can't cancel out your units, okay? So I know I've been telling you, cancel, cancel, cancel, and I'll give you the units, but when you're doing equilibrium constants, you've got to watch out, okay? Only the concentration of gases and substances in solution are shown, okay? So the concentration for pure liquids and solids are not shown. So you only do equilibrium constants for gases and things in aqueous solution, okay? It's because their concentrations change. Pure liquids are, like, essentially, like, I don't know, 100 molar or something, and you're not going to change that concentration very easily. Essentially, yeah, that's where you're taking H2O out, yeah. Okay, so if we wanted to write the equilibrium constant for this reaction, you see N2 plus 3H2 goes to 2NH3. Notice they're all gases, so we can put them in the equilibrium constant. We'll get equilibrium constant equals the molar concentration of ammonia squared. Why? Because we've got a coefficient there over nitrogen over hydrogen cubed, okay? Why is it cubed? Because of the 3 in front. Okay, I'll let you do these other two on your own. Notice this solid here, should that be in the equilibrium constant? No, but the aqueous, right, those should be, right? What about this liquid here? No, good job, right? All the gases, all the gases, yes. Okay, so we'll stop there, and we'll get the rest of the time to take the quiz.