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Studying for a test? Prepare with these 4 lessons on Kinetics.
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Video transcript
- [Voiceover] When you're studying chemistry you'll often see reactions, in fact you always see reactions. For example if you have hydrogen gas it's a diatomic molecule, 'cause hydrogen bonds with itself in the gassy state, plus iodine gas, I2, that's also in the gassy state, it's very easy to just sort of, oh you know, if you put 'em together they're going to react and form the product, if you have two moles of, hydrogen, two moles of iodine, so it's gonna form two moles of hydrogen iodide. That's all nice and neat and it makes it seem like it's a very clean thing that happens without much fuss. But we know that that isn't the reality and we also know that this doesn't happen just instantly, it's not like you can just take some hydrogen, put it with some iodine, and it just magically turns into hydrogen iodide. That there's some process going on, that these gaseous state particles are bouncing around, and somehow they must bounce into each other and break bonds that they were in before, and form new ones, and that's what we're going to study now. This whole study of how the reaction progresses, and the rates of the reactions is called kinetics. Which is a very fancy word, but you're probably familiar with it because we've talked a lot about kinetic energy. Kinetics. Which is just the study of the rate of reactions. How fast do they happen, and how do they happen? So let's just in our minds, come up with a intuitive way that hydrogen and iodine can combine. So let's think about what hydrogen looks like. So if we get our periodic table out, hydrogen's got one valence electron so if they have two hydrogen atoms they can share them with each other. And then iodine, iodine has seven valence electrons, so if they each share one they get complete as well. So let's just review that right now. So hydrogen this hydrogen might have one, well, will have one electron out there. And then you can have another hydrogen that has another electron out there, and then if they form a bond they share this, this hydrogen can pretend like he has this electron, this hydrogen can pretend like she has that electron, and then they're happy. They both feel like they've completed their 1S shell. Same thing on the iodine side. Where you have two iodines, they both have seven valence electrons. They're halogens, you know that already. Halogens are the group seven elements, so they have seven electrons this guy's got one here, this guy's got one here, if this guy can pretend like he's got that electron, he's happy, he has eight valence electrons. If this guy can pretend like he's got that one, same thing. So there's a bond right here, and this is why hydrogen is a diatomic molecular gas, and this is why iodine is the same. Now, when they're in the gaseous state, you have a bunch of these things that are moving around bumping into each other, I'll do it like this. So the hydrogen might look something like this, the hydrogen is these two atomic spheres that are bonded together, these electrons in between that are keeping them bonded. The iodine might look something like this, it's a much bigger molecule. Where it's bonded together like this, it's also sharing some electrons in a covalent bond, everything's probabilistic. So in order for these two molecules to turn to this, somehow these bonds have to be broken, and new bonds have to be formed. And what has to happen is that these guys, there's a ton of these guys. I could draw a bunch of them. Or I could copy and paste. So there's a bunch of... There's a bunch of hydrogen molecules around, and some of these iodine gas molecules around. So what has to happen in order for us to get the hydrogen iodide is, they have to collide. And they have to collide in exactly the right way. So let's say this guy, wish I could show it. Let's say he's moving, this is neat, I'm just dragging and dropping. But he's moving. He has to hit this hydrogen molecule just right, And maybe just right, if he just happens to hit it and bounce it with enough energy, then all of the sudden, let's say we get to this point right here, these electrons are gonna say "Hey, you know, "it's nice to be shared this way, "we're in a stable configuration, "we're filling the 1S shell, but look at this, "there's this iodine that's close by "and they really want me, "they're much more electro-negative "than me, the hydrogen". So maybe they're kind of attracted here, they don't know whether they wanna be here between that hydrogen and this right here between that. And so they kind of enter this higher energy state. And similarly, you know these guys they say, "Hey, wouldn't it be nicer, "I don't have to be here, I could kind of go back home "to my home atom if this guy comes in here". Because then we're gonna have, then we're gonna have eight valence electrons and the same thing's happening here. And this complex right here, this kind of, right when the collision happens, this is actually a state, this is the high energy state of the transition state of the reaction, and this is called an activated complex. Activated complex. Sometimes you know I just drew it kind of visually, but you could draw it like this. So hydrogen has a covalent bond with another hydrogen, and then here comes along some iodine that has a covalent bond with some other iodine,, but all the sudden these guys like to bond as well. So they start forming, so there is kind of a you know, there's a little bit of an attraction on that side too. So this is another way of drawing the activation complex. But this is a high energy state, 'cause in order for the electrons, the way you can think of it, to kind of go from that bond to this bond, or this bond to that bond, or to go back, they have to enter into a higher energy state. A less stable energy state, than they were before. But they do that if there's enough energy, 'cause you can go from, so you're going from both of these things separate, let me just draw them separate, so you have both of them separate, you have the hydrogen separate, plus the iodine separate. They go to this, which is a higher energy state. But if they can get to that higher energy state, if there's enough energy for the collision and they have enough kinetic energy when they hit in the right orientation, then, from this activated complex or this higher energy state, it will then go to the lowest energy state, and the lowest energy state is the hydrogen iodide. I wanna draw the iodide, and then the hydrogen. This is actually, this is actually a lower energy state than this. But in order to get here you have to go through a higher energy state. And I could do that with an energy diagram. So if we say that, let's say the X-axis is the progression of the reaction, and actually you know we don't know how fast it's progressing, but this you can kind of view it as time on some dimension, and let's say this is the potential energy. I wanna draw thicker lines. See this is the potential energy. Right there, let me make this line thicker as well. So this is the potential energy. So initially, you are at this reality, and we can kind of view it as the combined potential energy, so this is where eventually, we start off here, and this is the H2 plus I2 And a lower potential energy is when we were in the hydrogen iodide. So this is the lower potential energy down here. Lower potential energy down here. This is the 2HI, right? But to get here, we have to enter this higher activation energy, where the electrons have to get, they have to have some energy to kind of be able to at least figure out what they wanna do with their lives. And so you have to add energy to the system, you don't always have to add it, but if it doesn't happen spontaneously you're gonna have to add some energy to the system to get to this activated state. RIght, so this is when we were at this thing right here. We're there. So some energy has to be in the system, and this energy, the difference between the energy we're at when we were just hydrogen molecules and iodine molecules, and the energy we have to get to to get this activated state, this distance right here this is the activation energy. If we're able to get to somehow put enough energy in the system, then this thing will happen, they'll collide with enough energy and bonds will be broken and reformed. Activation energy. Sometimes it's written as Ea energy of activation, and in the future we'll maybe do reactions where we actually measure the activation energy. But the important thing is to conceptually understand that it's there, that things just don't spontaneously go from here to here. And I won't go in deeply into catalysts right now, but you've probably heard of the word catalyst or something being catalyzed. And that's something, some other agent, some other thing in the reactions. So right now, so right now we're doing, we have H2 plus I2, yielding 2H hydrogen iodides. Now you could have a catalyst, and I'll just say plus C. And I actually don't know what a good catalyst would be for this reaction, and how a catalyst operates is, it can actually operate in many many different ways, so that's why I don't wanna do it in this video. But what a catalyst is, is something that doesn't change. It doesn't get consumed in the reaction. The catalyst was there before the reaction, the catalyst is there after the reaction. But what it does is it makes the reaction happen either faster, or it lowers the amount of energy for the reaction to happen. Which is kind of the same thing. So if you have a catalyst, then this activation energy will be lower. And what it does is, it makes it, it might easily, it might be some molecule that allows some other transition state that has less of a potential energy so that you require less heat or less concentration of the molecules for them to bump into each other in the right direction, to get to that other state. So you require less energy. So given how we understand how these kinetics occur, these molecules interact with each other, what do you think are the things that will drive whether a reaction happens or not? I mean we already know that if we have a positive catalyst, there's something called a negative catalyst that will actually slow down a reaction. But if we have a positive catalyst, it lowers... Obviously it lowers the activation energy, so this makes reaction faster. More molecules are gonna bump into each other just right to be able to get over this hump 'cause the hump will be lower, when you have a catalyst. Also if you increase the concentration. Right? If you increase your concentration, of molecules, if the concentration goes up, then you just have more stuff to bump into each other. Right? There's just the likelihood. Everything is probabilistic. You know when people write these reaction equations, it all seems nice and simple and very clear and it happens. But no, in the real world, you just have things bumping into each other. And when we do biology videos it'll be fascinating to talk about. Because all of, every biological process is really just a chemical process. And it's really just a byproduct of all of these things bumping into each other. And you can imagine the more concentration you have of the things that need to bump into each other, the more likely you're gonna get just that perfect bump, and that perfect amount of kinetic energy for the reaction to happen. And actually I'll make a little other note here. This reaction you might say, "Okay I have some, "let's say I'm at this energy, "how do I ever get over this? "How does this ever react?". Well remember, in a gas, the kinetic energies of all of the molecules, they're not uniform. Some gases have, some molecules will have higher kinetic energy, some will have lower, temperature just gives you the average. So there's always some probability that two maybe high kinetic energy molecules will bump into each other just perfectly, surpass the kinetic-- So they have enough kinetic energy to get into the activation state, and then they can go to the lower state, which is the hydrogen iodide. So even at a, at all temperatures this will occur, but obviously if you increase the temperature, if you increase the temperature that reaction is more likely. So that's the other one. So temperature. Temperature is actually probably the biggest. Temperature is probably the single biggest thing that will make the reaction happen faster. So all of these things, you want higher temperature, higher reaction. And then if you just wanna think about the molecules itself, if you have molecules where their original bonds are weak, they're more likely to be able to interact. And there's other things you could talk about, the molecular shape, how available certain atoms are to interact with other atoms, and that really becomes significant when we start going into biology. And then the last one, and you probably realize this, is just the surface area. If you increase the surface area, so we were just doing gas-gas interactions, which almost by definition have pretty good surface area interactions. But if the surface area goes up, then the reaction also goes up. The reaction rate. And how do you think about that? Well, think about the reaction of, think about the reaction of, you know we've done this multiple times. Sodium chloride, solid, so solid salt, plus liquid water, leads to sodium, well we could think of it a lot of different ways, but we could think of it as sodium ion aqueous, plus chloride anions, there's a cation anion, aqueous. So it gets dissolved. And how does that happen? If you have a big block of ice, not of ice, of salt. You have a big, I'll do salt in grey. If you have a big block of salt in there, and you have, so there's a bunch of sodium and chloride atoms in it. And you have water all around it, the water is only gonna be able to interact with the surface molecules, and slowly dissolve away the salt. Slowly make polar bonds. These are actually, well, polar dipole bonds with a different, the different sodium or chloride ions. But if you were to break this up into smaller cubes, if you were to break it up or really crush it into really small pieces, then all of the sudden the surface area that the water molecules can interact with, it can actually interact with more of the sodium chloride. So the reaction will happen faster. So surface area, you increase the surface area of interaction, and you'll also increase the reaction rate. If you're trying to do it with two fluids, what you could do is you could kind of spray one fluid into the other, so you have little droplets so you also increase the surface area. So anyway, this is kind of an introduction to the idea of kinetics. But hopefully gives you a sense that these reactions, and I want you to really think about chemistry this way, not think about it as, "Oh it's just some formula I have to remember". That that these really are bumps and bruises between atoms, that it's probabilistic and it's messy and we really have to think about what will make it more likely that these things collide in just the perfect way for the reactions to happen.