tinywow_video_to_text_65331227.txt

So now in this next section, we're going to do a little bit of a chemistry review on thermodynamics and kinetics, and really use it as an introduction into enzyme kinetics. So show it with a little bit of a view of thermodynamics, and both the topics here on thermodynamics and kinetics will definite...

So now in this next section, we're going to do a little bit of a chemistry review on thermodynamics and kinetics, and really use it as an introduction into enzyme kinetics. So show it with a little bit of a view of thermodynamics, and both the topics here on thermodynamics and kinetics will definitely cover in much more detail in the general chemistry section. So a little more review and a little more of just setting the foundation for enzyme kinetics here. So first thing I'm going to talk about is Gibbs free energy, and that's delta G. So this is the big thermodynamic quantity that tells us whether or not a reaction is going to happen or not. The word we use is whether or not a reaction is spontaneous or not. And so in this case you want to know that when delta G is negative, and you should definitely know that delta G being less than zero, that just means that delta G is negative. So that's when a reaction is spontaneous, we call that reaction exergonic. In other words, for the less commonly used, but in other words you probably should file away. So whether you say it's spontaneous or exergonic, same diff. Delta G is negative. It means the reaction is going to proceed downhill in terms of free energy. So you should know that Gibbs free energy here is the energy available to do work. And when it decreases, it means you've used up, or at least some system has used up its energy to do some work. And when the system uses up its energy, that's a spontaneous reaction. It's using up its own energy. So for non-spontaneous reactions, when delta G is positive, so these are up hill. They require the input of energy from outside of the system. So from the surroundings and stuff like this, we call them endergonic or non-spontaneous. Again, both words you should associate with delta G being positive. And then finally, a reaction is going to proceed in the spontaneous direction until it reaches equilibrium. And when it gets there, that's when delta G is going to equal zero. So it turns out a reaction is going to be spontaneous or non-spontaneous under one set of conditions. And that's going to tell you, based on this condition, whether it's going to run in the forward direction or the reverse direction, whichever one is spontaneous. So and it'll run as it does, you're going to find that delta G value is going to approach zero over time. And when it finally reaches equilibrium, that's when it has reached zero. Now, because we can use delta G as an indication of when we've reached equilibrium, so it makes sense to also talk about the relationship between delta G standard and this little circle right here means under standard conditions. And those standard conditions mean, you know, all reactant concentrations are one molar or if you have gases, they all have partial pressures and one atmosphere. And technically it doesn't mean 25 degrees Celsius, but most of the time if we're talking about standard conditions, you're probably going to be talking about standard conditions at 25 degrees Celsius. So I'm going to kind of throw it in there. 25 degrees Celsius, 298 Kelvin, same diff, though technically not a part of standard conditions. All right. But when you have delta G under standard conditions, there's a relationship between it and the equilibrium constant K. And just keep in mind that, you know, what the equilibrium constant is. If you've got a reaction like A going to B, the equilibrium constant is going to be the ratio of the concentration of B over the concentration of A. So ratio products to reactants. So and if we had coefficients, it'd be the coefficients that we show up as exponents. And again, more detail on this in the gen chem section. So the big thing here is that when you've got a very, very large equilibrium constant, a very large value for K, well that means your numerator is big and your denominator is small. You've got a lot of products, not a lot of reactants. And so when an equilibrium constant is a lot bigger than one, usually think like bigger than 1,000, that's a reaction that's going to favor the products. There's going to be a lot more products than reactants at equilibrium. So on the other side though, what if your equilibrium constant is a really small number? Now it's never going to be a negative number because concentrations are never negative. They can't have less than none. So, but if it's way smaller than one, so like 0.001 or 1 times 10 to the negative 8 or something along these lines, when you've got a really small equilibrium constant, what that means now your denominator is bigger than the numerator by a fair amount. So you have a lot of reactants at equilibrium and not very many products. And so in that case, that's why we call that a reaction that favors the reactants. And then finally, if you've got an equilibrium constant somewhere in the ballpark of one, so somewhat close to one either, a little bit higher, a little bit lower. So that's going to be a reaction that favors a fair amount of both products and reactants at equilibrium. Now, it's less convenient to see this equilibrium constant because you look at one number and it's going to tell you something about that reaction, whether it favors the reactants of products or both at equilibrium. Well, it turns out you can relate delta G standard to that equilibrium constant as well. And it turns out it's through this lovely equation here, delta G standard equals negative RTL and K. Now, the truth is, I don't really care about this equation. You're never going to do calculations with this equation, not even in the Gen Chem section because it has a natural log in it and you don't get a calculator. So that's not the big deal. The big deal though is is the equation that describes the relationship between delta G standard and the equilibrium constant. It's the way this works. If delta G standard is less than 0, I E negative, well that means that that reaction is spontaneous under standard conditions. Well, then if it's spontaneous under those standard conditions where you have one more concentrations of all reactants and products, well, if it's spontaneous in the four direction, well, then by the time you get to equilibrium, you're going to have way more products than reactants and you're going to end up with your equilibrium constant being larger than 1. And it turns out there's mathematically, you can work that out, K is larger than 1 than Ln of K, well, Ln of a number larger than 1 is positive. And when you factor that negative sign in, you can have delta G standard, that is negative. So delta G standard being negative means your equilibrium constant is going to be larger than 1, favoring the products. And it kind of makes sense, right? So we associate delta G being negative as being spontaneous. Well, if it's spontaneous under standard conditions, then you should end up with most of the products in K's negative 1. We've got the opposite as well. If delta G standard is positive, well, now it's not the forward reaction that's spontaneous. That's non spontaneous. And so once you find the reach equilibrium, you're actually going to end with more reactants than products. So if you have more reactants than products, then all of a sudden, this ratio is going to be smaller than 1 instead. So it makes sense, again, with delta G standard being positive, it's the reverse reaction that's spontaneous. The forward one is non spontaneous. So you end up with more reactants. And then finally, if your delta G standard was 0, that is the one instance where you end up with an equilibrium constant of 1. So this is, you know, very rare, not the most common thing in the world. It could happen, but much more than the time we're probably just going to be talking about when delta G standards either negative, and K is bigger than 1 or when delta G standard is positive, K is less than 1. So now we're going to talk about catalysts a little bit, and this really is just so we get an introduction into enzymes, so enzymes and kinetics. So enzymes are biological catalysts. We just want to talk about catalysts ever so briefly. So what a catalyst is ultimately going to do. So to convert reactants into products, you've got to go over some energy barrier. We call that energy barrier the activation energy. So the top of the hill here, the top of that energy barrier, you've got your transition state or activated complex. And again, also something will go into more detail in the general chemistry section of course. So what a catalyst is going to do is it's going to provide an alternate pathway mechanism with a lower activation energy. And if you have a lower activation energy, then the reaction is going to proceed over that now lower energy barrier much more quickly. So the catalyst that you need to know, again, will also get in cover of this in Gen-Chem. So it speeds up a reaction and you should realize that it's going to speed it up in both directions. It's not just lower in the forward direction here. It's also, if you were going backwards, it's lower in that direction as well. And so it speeds up both the forward and reverse directions. So for the reaction. And it lowers the activation energy in both directions, both the forward and reverse reactions. So and it does this by providing some sort of alternate pathway or mechanism for the reaction to occur. Now one thing you should notice, a catalyst is not part of the overall balanced chemical reaction and therefore it's not produced or consumed net in the reaction. So overall it's not consumed in the reaction. So one catalyst will catalyze a reaction and then it'll, you know, interact with some more reactions and catalyze it again. Same thing with enzymes in your body. They're a biological catalyst and one enzyme can catalyze the same reaction over and over and over again. So then finally, because it's speeding up the forward and reverse direction, reactions at the same time, it does not shift the equilibrium in the reaction. So if you've got a reaction that's already reached equilibrium and you've got a catalyst, it will do nothing. A catalyst is not going to shift equilibrium towards the reactants or the products at all. You're just going to get to equilibrium faster. So if you're already at equilibrium, adding a catalyst will accomplish nothing.