 Okay guys, let's finish up our talk about system and surroundings. If you guys haven't noticed, the test grades have been posted. They're all up on Blackboard. I'll be handing them back today. So we'll do it close to the end of class. But anyways, recall, in any experiment, you've got two different places in the universe, right? So the universe is composed of two different areas, right? The surroundings and the system. The system is whatever you're working on. The surrounding is everything else in the universe. Notice you've got a system here that's giving off heat into the surroundings. So what they're showing here is the temperature of the surroundings is increasing. This would be like something if you, a reaction that was occurring and you put your hand on the beaker and it was, your hand was getting warm, okay? In this case, the opposite is happening. The heat is flowing into the system. In this type of a reaction, you would put your hand on the beaker and your hand would be getting cold, okay? So energy change usually comes in the form of heat like is shown here. And fortunately enough for us, this change can be measured and that's what we're going to learn about. So the first law of thermodynamics. The first law of thermodynamics states that the energy of the universe is constant. So you might have heard this as energy can be neither created nor destroyed, okay? So what you want to realize is that the total energy of the universe here equals the energy of the system plus the energy of the surroundings, okay? So if we set the energy of the universe equal to zero, which is okay to do, what we get out of this equation is this new equation, the change in the system, change in the energy of the system equals the negative change of the energy of the surroundings, okay? So this makes sense. However much heat is flowing out of your beaker, that's how much heat the surroundings are absorbing, okay? That's all that's saying. Okay, so let's talk about spontaneous processes and we'll talk about them in relation to chemical reactions. This is a pretty good simple diagram that kind of indicates what kind of reactions you might see actually. This is a macromolecular process that kind of resembles what reactions look like, okay? So spontaneous processes are processes that take place naturally with no apparent cause or stimulus, right? So you can imagine this being a spontaneous process, maybe the man doesn't even have to be there, the rock would just fall off, okay? Fall, roll down, okay? So that's what they're saying here. Notice here, this would be a spontaneous process if the rock was right here. But so once this guy pushes it up to about right there, then it's a spontaneous process. So that's what this is saying. It's a spontaneous process once started. So you need some little bit of energy to push it up to that higher position. We're going to call that an activation energy, that little push of energy. And then you can see this guy here like Sisyphus, right? Non-spontaneous process, always pushing the boulder up the hill, okay? So it takes energy to go this way, right? External energy source. But it takes no external energy source to go this way. Does that make sense? It takes a little bit of external energy to push that guy up, okay? So let's, this thing will say, let's talk about these reaction spontaneity and free energy. So exergonic processes are processes that give up energy as they take place. So that means the surroundings are becoming warmer, okay? So energy or heat is flowing out of like your reaction vessel, if you will. Okay? So energy can appear if you want to think of energy as a reactant or a product, right? Energy in this case is a product. It's coming out of this reaction. And we can write it on the right side of the reaction just like we would write the products of the reaction. Remember, kilojoules or kilocals like is shown here or calories, those are all energy units. So this would be an exergonic process where we have the energy on the right side of the equation or your reaction vessel giving off heat. Endergonic processes, on the other hand, are processes that gain or absorb energy as they take place. So this process here, this mercury-2 oxide going to its elements, mercury and oxygen, it takes energy, you need to pump energy into the system to allow it to react. Okay? So in this case, you're going to put the energy on the left side of the equation. So does this make sense? This reaction wouldn't occur unless I put energy into it. That's why I'm putting it as a reactant, okay? Because normally, if you have A plus B here, goes to C, if you don't put B in that chemical equation, it won't go to C, okay? So you can think of this just as another reactant or as another product depending on what's going on. Okay? So let's consider that general reaction. AB plus CD goes to AD plus CB. So this kind of double displacement reaction, if you will. So what do we have here? We've got A bonded to B, right? So that bond is actually holding those two atoms together. In fact, it's stored chemical energy, okay? So it's energy itself. And what happens when a reaction occurs is that bonds must break and bonds must form, okay? Breaking bonds requires energy. Remember we talked about the activation energy. You have to push the boulder a little bit up the hill before it will fall down automatically, okay? This activation energy, you need to kind of push these reactions to get them going. Okay? So for an exergonic reaction, we can think of a general description of the reaction like this, right? So if we look at this energy diagram, we're going to have our reactants at a higher energy level than our products. Why is that? So this is energy, and this is time, right? So you start the reaction and the reaction's finished. Remember, exergonic things release energy, right? So what happens to the energy from here to here, right? That all gets released as heat, okay? So that's going to be coming out of the reaction like that. Does that make sense? Hopefully that makes sense. I know it's Friday morning at 9 o'clock. Okay, so this is the reaction that releases energy. And if the energy required to break the bonds is less than the energy released when the bonds are formed, there's a net release of energy. So what is this saying? That forming these bonds doesn't take as much energy as forming these bonds. They're very highly energetic bonds, and these aren't as energetic. That's why the thing goes down. The reaction goes down in energy. So what is this saying? The reactant bonds must be broken in the reaction, requiring energy, and the product bonds are formed, releasing that energy. So an endergonic reaction is just the opposite of that. So if the energy required to break the bonds is more than the energy released when the bonds are formed, then you need to have an external energy source, like this guy pushing the boulder up the hill. Okay, you need to have the band there pushing it up. In this case, we'll have energy as a reactant in these equations. So if you recall a combustion reaction, this is a good example of having a reaction that is exergonic in nature. It's going to lose energy as heat. Probably you've been near a campfire or something like that and felt the heat coming off of that campfire. That's due to the fact that it's an exergonic reaction and it's releasing heat or releasing energy in the form of heat. The decomposition of ammonia would be an example of an endergonic reaction. You need 22 k-cals of energy to pump into this system to decompose ammonia to its constituent elements, nitrogen and hydrogen gases. Okay, so let's talk about something that's very similar, enthalpy. Enthalpy is the term that we used for energy gained or energy lost by heat. Okay, and the symbol for enthalpy is H. What we're going to be noticing or using a lot more of, though, is the change in enthalpy because we can't really measure the inherent enthalpy that these things have, but we can measure the change going from reactants to products. Okay, so the change in enthalpy will be the actual amount of heat released or needed by the reaction. So the change in enthalpy, another way of stating that, is the energy difference between the products and the reactants of a chemical reaction. So when energy is released as heat, this is known as an exothermic reaction. So we say that that enthalpy change is negative. So in this case, what's happening? Our initial energy starting point is here. Energy level one is there. Energy level two is here. Is energy two lower than one? It's lower than one. So if we took E2, subtracted E1 from it, this is going to give us the change in enthalpy. That'll be a negative number if E2 is lower than E1. That's an exothermic reaction. Exothermic reactions are almost always spontaneous processes. We just need that little activation energy to push them over the hump. If energy is absorbed, we call it an endothermic reaction and the enthalpy change is positive. That would happen in the opposite situation if we had an endothermic reaction. The products would be at a lower energy level than the reactants. So in order to get up that hill, we're going to have to add heat to this reaction. If we say that this is E1 and this is E2, E2 minus E1, is that going to be a positive or a negative number? Positive number because E2 is bigger than E1. So that's going to be delta H. It's going to be positive there. It's going to be negative there. This is an exothermic reaction. This is an endothermic. Notice where heat is also in these reactions. Here it's a product. My hand is getting hot if I'm touching the beaker. Here it's a reactant. My hand is getting cold if I'm touching the beaker. Remember, a spontaneous reaction occurs without any external energy input. Most, but not all exothermic reactions are spontaneous. Thermodynamics, which is what we're essentially studying right now, is used to help predict whether the reaction will occur. We're using just this term enthalpy to predict whether these reactions will occur or not. Unfortunately, I guess for general chemistry students, there's another term that has to be added into the enthalpy term, and this is called entropy. You may have heard of entropy before. I don't know, there's a common joke that you might see that's like entropy works or things like that, and that's always placed on a room that has a lot of disorder in it. Look at these. This should say delta S, here. Delta S, delta S, instead of delta H. When delta S is positive, what is delta S? Delta S is entropy. When delta S is positive, let's look at these two pictures. Do you think that would be a spontaneous process? If we have gas in here and I open this stopcock, do you think the gas would flow into there spontaneously? Probably, right? Probably. What about this process? This thing, you open the stopcock. Do you think all of the gas particles will flow into there on the left side? No. So cool. You guys already understand entropy. Entropy states that the more disorganized the system, the more favorable it'll be. This happens to me like every week. I clean up my bedroom on the weekend, and then by the end of the week, it's all messy. Why is that? Because it takes energy to keep it maintained. Things will get disorganized inherently. So it takes energy for me to clean up my bedroom, but it doesn't take much energy for it to get disassembled again. So this is like entropy. Right here, this doesn't take any entropy to take these gas molecules and to move them over there. It just takes the activation energy of us opening that stopcock there. But here you can imagine it might take a lot of energy to try to get all those gas particles into one side of those two bulbs like that. And then here I guess we got some exothermic endothermic questions. Which is this process? Is this exothermic or an endothermic process? If our delta H is negative 16, exothermic. And this one here, if it's positive, endothermic. So we've been talking about with this entropy thing, the second law of thermodynamics. So you already inherently knew it though. An isolated system will tend to increase in entropy over time, get more disordered over time. So the universe spontaneously tends towards increasing disorder or randomness. Like here, if we're looking at these two jars of jelly beans, which one do you think would have the more entropy on the left or on the right? The right, right? That's more disorganized. Took energy to make it look like this. Didn't took very much energy to make it look like this. So again, entropy is S. We're going to usually be talking about delta S, which is the change in entropy. Entropy itself is a measure of randomness or disorder of a chemical system, in this case when we're referring to chemistry. So if I have high entropy, we say this is a highly disordered system. Like here, absence of any regular repeating pattern. Low entropy is a well organized system, such as this or crystal structures or complex molecules, things like that. So why biological systems are so interesting, right? Because they go against the second law of thermodynamics, right? So it takes a lot of energy to, you know, make a person through chemical reactions. Like all the chemical reactions that are happening in our body right now are going against the second law of thermodynamics. On the surface, it seems that way, you know? But in actuality, since the universe has to tend towards increasing entropy, what we do in our bodies is couple reactions and have very highly favorable reactions. What we say couple or react together, make each other react with highly disfavorable reactions. That's why we give off heat, too, right? So we can feel all the exothermic processes going on inside of us, like when you hold an ice cube. That's why it melts or something like that, you know? Or when you're next to somebody and you're like, get away from me, it's too hot or you're a stuffy place, you know? So let's talk about the change in entropy of a reaction. Same thing like the change in enthalpy is the entropy of the products minus the entropy of the reactants. So like what we said, a positive change in entropy means an increase in disorder for the reaction. So melting, going from a solid to a liquid, we would predict that would be a positive or negative entropy change. Positive, right? Because why is that? Because it's getting more disordered, right? More disordered. What about going from a liquid to a gas? That a positive or negative entropy? Positive as well, positive as well, right? So positive, remember, for entropy is good. Negative for enthalpy is good, okay? Or good as in favorable for making the reaction go. I don't know if there's inherently good or bad in chemistry. And dissolving is also an entropically favorable process. Going from this highly ordered system with the solvent being all around the solute to the disordered system of the solute breaking up into its ions and flowing around. So again, negative entropy means decrease in disorder. So that means increase in order, right? For the reaction. So what do you think? Which substance has greater energy? Helium gas or sodium solid? Helium gas, why would you say that? Because it's a gas, right? Because solids are like what? What are the particles and solids like? What are the particles and solids like? All stuck together, right? And gaseous particles, are they stuck together? No, they barely even see each other, right? What about liquid water or gaseous water? Gas water, right? Or steam, right? That stuff would be much more entropically favored. Okay, so we've been talking about reactions, spontaneously going, non-spontaneous, all this stuff. Well, why would a reaction not spontaneously go? Why is that? Well, it's because we have substances that are stable, okay? So substances that are at a low energy state relative to their makeup of chemical bonds will not undergo spontaneous changes. They're happy, if you will, if you want to anthropomorphize molecules. They're happy if you want to make them to have human emotions, right? Molecules. They're happy at the state that they're at, the low energy state. So why is like, you know, why would you want to walk up this huge hill if you could be just chilling at home in your air conditioner, right? You're not going to do it because it takes a lot of energy. Why would you do it? What about going downhill, right? Flying downhill, yeah, that's what you want to do, right? You want to get back down to your home, right? But you can see here is bromine, solid, liquid gas, increasing in entropy, increasing in entropy. But stable substances are substances that don't undergo spontaneous changes at the prevailing conditions, temperature, pressure, etc. So let's talk about these two things together, okay? So now with these two new concepts of enthalpy and entropy, you should be able to predict whether a reaction will occur spontaneously or not. So you would imagine that if it was a negative enthalpy and a positive entropy that this reaction was displaying, that that would always go, right? Because positive is good for entropy and negative is good for enthalpy. So that's always going to be a spontaneous process. And that's what this says here. A process will always be spontaneous if it gives up energy, right? Gives up energy, what does that mean? Negative delta H, right? Negative delta H giving up energy. And the entropy of the system involved in the process increases. So what would that mean? Positive delta S, yeah, very good, positive delta S. Okay, does that make sense? Okay, cool. So a process that absorbs energy will be spontaneous only if an increase in entropy of that system occurs, but that is large enough to compensate for the increase in energy. Okay, so what you'll find is that the entropy term is very, very small relative to the enthalpy. Okay, so it's not, if you've got a big change in enthalpy, you're not usually going to have your entropy change be able to overcome that enthalpy change. Okay, so a process that causes an entropy decrease in the system will be spontaneous only if the decrease in energy also occurs that is large enough to compensate for the entropy decrease. This happens more often than this one. Okay, so in other words, if you're exothermic, you have a negative delta H and a positive delta S, you're always spontaneous. If you're endothermic, positive delta H, and have a negative delta S, you're always non-spontaneous. Okay, for any other situation, it actually depends on the relative size of delta H and delta S. Okay, this is what the previous slide tells us just more concisely. Okay, so how do we figure out if these changes are big enough or small enough? Well, lucky for us there's another equation that'll tell us exactly that, right? And it's this equation, delta G. Okay, so delta G stands for Gibbs free energy. We could just call it free energy because Gibbs, the guy Gibbs, that this is named for is long since dead, so I don't think he'll have any problem with it. But that's why it's this big G here. So this represents the combined contribution of the enthalpy and entropy values for a chemical reaction. So remember when we said exergonic and endergonic, right? Delta G, if you've got a positive delta G, that's when you're endergonic. When you have a negative delta G, that's when you're exergonic. So you've got to go back to the initial part of this discussion to connect those two. Okay, so in other words, if you've got a negative delta G, you're always spontaneous. If you have a positive delta G, you're never spontaneous. And how do you get delta G? You guys can't read it. Delta G equals delta H minus T. T is temperature, okay? And Kelvin times delta S. Temperature is in Kelvin. Those little kind of degree signs, do you guys see that delta G with the little, those things? That just means at its standard state, okay? 273 Kelvin, 1 atm, okay? Atmospheric pressure, 0 degrees Celsius. Okay, here we go again. Let's put it all together. This is like our bookkeeping slide. So, of course, we're going to need to know both the delta H and the delta S in order to predict the sign of delta G. We'll also need to know what the temperature is of the reaction that we're analyzing. So, if we have a positive delta H and a negative delta S, delta G will always be positive. That means it's going to be non-spontaneous, regardless of temperature. Okay? If we got delta H is negative, delta S is positive, delta G is always negative, regardless of temperature. And if delta S is positive and delta H is positive, then we need to know what the temperature is to determine what the sign of delta G is and the same is true for when they're both negative. We need to know the temperature. So, we're just going to have to plug all these values in. You're going to be able to get these values from the problem, okay? So, that was all the thermodynamic stuff. Okay, let's talk about some kinetics. So, until now, we've been talking about reactions like A plus B goes to C. Okay, just like, here's products, here's reactants. Okay, this is like talking about, I don't know, race or something and saying like a horse race and saying or a person race, right? That here are all these people and here's the person who finished first and here's the person who finished second. Here's the person who finished third. How did that happen, right? How did the person who finished first trip the person who finished second or did the person who finished last run the wrong way or something like that? How do we know what happened during the race? We don't know. So, this is what we've been thinking of is kind of like the box score, if you will, of the reaction, right? It just tells us who played the game and what the outcome of the game is. And we knew now the delta H of the reaction. So, we never considered, well, how fast is this reaction going? What are the reactant and product concentrations when the reaction is complete? And will this reaction be spontaneous? So, you can see, if we take measurements at different in time intervals, you can see the relative concentrations of the reactants decreasing and the relative concentrations of the products increasing, right? So, you would expect this. Can I erase all this stuff on the board? You would expect this, but you probably just never thought about it before. Draw another graph up here. Okay, so, if this is concentration and this is time, right? Initially, we would have the product concentration. Remember, brackets around something is molar units, good concentration. Very high, right? And after a while, we would expect that to go down, down, down, down, down, probably to zero, like that, right? And if we looked at the reactants concentration, or the products, oops, alright, that's the reactants. The products, concentration, alright, that makes sense. The other didn't, right? We would expect there would be no products at the beginning, right? No products at the beginning. And then after time, they would increase, increase, increase until like that, right? This would be this reaction, R goes to P. Like that. Does that make sense? If we look at that graph. Okay, so if we took the, if we took a snapshot of this reaction at this, at time here, right? We wouldn't, we wouldn't get a picture of what happened over here, but we'd get kind of an idea of what's going on in between, right? Now we could take a snapshot here, and we could see what happened. Well, the product concentration increased to there, and the reaction, reactant concentration decreased to there, right? Now if we take it here, then we could see that. If we take it here, we could see that, that, that. Okay, so we can see what's the rate of the product concentration increasing, and what's the rate of the reactant concentration decreasing. So, what we're talking about now is like I said, chemical kinetics. Kinetics just means the study of how things are the rates of things. Okay, so you can have kinetics of like a race, you know? Reaction mechanism, this is the step-wise changes that reactants undergo in their conversion to products. Okay, so maybe R doesn't go directly to P. Maybe R goes to some other thing first. Maybe it goes, R goes to A, right? And then A goes to C, and then C goes to P, right? But if we're only looking at the beginning and the end of the reaction, we wouldn't see any of this intermediate stuff, okay? So the reaction mechanism would be R to A to C to P, okay? And in fact, you can, you can predict like the way that these things are going to react by, well, you can predict the reaction mechanism by the structures of these various molecules that you're using to react. Okay, so there's a couple other definitions. Well, thermodynamics we already went over. It's the conversion of chemical energy to heat and the reverse of that, okay? That's what we talked about for the first part of lecture today. An equilibrium is the dynamic balance between forward and reverse reactions and how to influence this balance by outside forces. So we've also only been talking about chemical reactions going like this from reactants to products and said, that's it. You know, that's what they do. But in actuality, all chemical reactions are reversible, okay? So not only do we have some concentration going from reactants to products, but we also have some products going back to reactants, okay? This happens in almost all chemical reactions. And that process of going back and forth is known as equilibrium, okay? Okay, so chemical kinetics is the study of changes in the concentrations of reactants and products as a function of time. Reactant concentrations decrease rapidly in a reaction with a high rate and slowly in a reaction with a low rate, okay? So you can imagine this would be a higher rate of reaction than maybe this reaction would be. Like that, okay? Yeah, and reactions proceed at a wide range of rates. Hopefully you can think of a couple of reactions that occur instantaneously, like the detonation of TNT, right? That's gonna occur like instantaneously, explode, right? But what about the process of aging, right? That takes a long time. What about the process of cooking a steak on the grill? That takes an intermediate amount of time. Or the process of the reactions that you do in chemistry labs. Although you might think it doesn't take, or it takes a long time, it's also an intermediate kind of reaction. So you can see here explosion, fruit, ripening, I guess, or molding or something, I don't know. And rusting, this would be another, like, long-term reaction, baby going to grandma, right? Long time. Okay? So let's just finish this, these couple of slides, and then we'll hand the test back. So the rate of anything, like we said, is a change in some variable over time, right? What's the rate of the stock market decrease or something like that, right? You can think of a rate of, I don't know, what rate would you think of? I heard about that on the radio this morning, okay? In this case, we're going to be talking about reaction rates, right? Reaction rates are the changes in concentration of reactants and products per unit time. So remember any rate, like a speed, is the change of position over the change in time, okay? So that would be like X2 minus X1 over T2 minus T1 or the change of X over the change of T, okay? But when we're dealing with reaction rates, we're dealing with two different rates, okay? So remember that the concentration of reactants is going to decrease over time, okay? So we're going to have a negative rate for reactants, okay? But when we're talking about products, those are going to increase over time, so they're going to have a positive rate, okay? So it's positive numbers going up, right? Negative numbers going down, okay? So you've got to watch that. Notice here you can see, you can see the rate of product decrease here in this particular instance, this reaction. There you can see product increase or product increase, reactant decrease. And, well, it's pretty much described to you right there. Moles per liter, molarity. Whenever you see those brackets, it's moles per liter. And this figure is demonstrating the change of purple reactant to green product over time from a molecular perspective. You can see what's going on. Here is some kinetic data being monitored by color, okay? You can see the reactant decreasing over time because it's getting, the whole solution's getting lighter and lighter and lighter, okay? So the rate of color change can aid in the calculation, calculating the rate of chemical reaction. When I was an undergraduate, I was doing research in biochemistry lab and we were doing this, like, doing reaction rates with enzymes that had these fluorescent tags on them and it would cleave the fluorescent tags so it would be like a decrease in, in this case, fluorescence over time, not color change, but it's the same thing, okay? So I'm going to leave you guys today with this calculation. I'd like you guys to try it on your own. So try to calculate the average rate of this reaction and you can read it on your own and that's how you do it. Okay, so I have, if you don't mind, don't leave yet before you pick up your test, okay? So your test and all of the other things that I have for you guys, so if you could wait. Just to let you know some details about the test. The high on the test was 100, I think four. Yeah, thank you guys. Four people got 100 on this test. There was many people more, well, a few more people that got 95 or better, okay? So this is a very good test. I was very impressed with the way you're working. Some people still need to probably work a little harder and I know everybody's working really hard these days, but of course if you're not getting it, you're going to have to work a little bit harder, okay?