Explore the role of thermodynamics and kinetics in chemical reactions. Understand what it means for a reaction to be thermodynamically and/or kinetically favored. Dive into the concept of free energy change, activation energy, and how enzymes can speed up reactions. A perfect blend of biochemistry and real-life analogies! Created by Jasmine Rana.
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- At3:26, shouldn't it be "transition state", rather than "intermediate"? I was under the impression that whenever there was an intermediate, it would be a local minimum.(39 votes)
- what is the differences between thermodynamics and kinetics?(16 votes)
- Thermodynamics, simply defined, is the branch of physics that deals with the conversion of different forms of energy, and the relations between heat and various energy forms such as mechanical, electrical, or chemical energy.
Kinetics deals more with the actions, or forces, that cause various motions (also known as dynamics). It is concerned with the rate of reactions.(31 votes)
- So I guess you can say an Airplane is the enzyme in your venture to getting to your special someone. Without the enzyme, you're travelling across the world with only a row boat.(15 votes)
- at1:13whats rxn stand for?(0 votes)
- Wouldn't the reaction on the left (the spontaneous one) be more kinetically favorable then the reaction on the right (non-spontaneous one)?
You couldn't have a reaction that is Kinetically favorable but thermodynamically unfavorable?(3 votes)
- You're correct about the reaction on the left being both spontaneous/thermodynamically favored (deltaG<0) AND kinetically favored (activation E)fwd < (activation E)reverse.
It is possible to have a product that is thermodynamically unfavored but kinetically favored. But I think if you're talking exclusively about a forward vs reverse reaction (like A<=>B), because the transition state is the same energetically for the forward and reverse, the thermodynamically favored reaction will always be kinetically favored. This is because the higher E/less stable reactant is energetically closer to the transition state than the lower E/more stable product.(4 votes)
- To clarify:
Forward Reaction: A -> B
- Thermodynamically favourable (G<0 then spontaneous)
- Kinetically favourable (Assuming Ea is provided from the environment or ignition source. Small Ea)
Backward Reactions: B -> A
- Thermodynamically unfavourable (G>0 then non-spontaneous)
- Kinetically unfavourable (Larger Ea)
If this reaction is an equilibrium reaction, how is the backwards reaction even possible? Sure, the forwards reaction will decrease or deaccelerate until the quotient reaction equal the equilibrium constant. But still, no B can be formed if it is thermodynamically and kinetically unfavourable.(4 votes)
- Are all thermodynamically favoured reactions spontaneous? That is, do all thermodynamically favoured reactions have reactants high in Gibbs and products low in Gibbs?(1 vote)
- Thermodynamically favored reactions do have a negative delta G (high free energy in reactants and lower free energy in the products) and thus can be called spontaneous! However, that doesn't mean that every spontaneous reaction will happen immediately!
Think about it like this--burning a log is thermodynamically favorable, however, you need an input of energy for the process to begin (IE: lighting a match). Logs don't randomly catch on fire! (thank goodness!)(4 votes)
- So does this mean that the reaction on the left is both thermodynamically favored (delta G is zero, thus spontaneous) AND kinetically favorable (activation energy is lower), and that the reaction on the right is both thermodynamically unfavored and kinetically unfavored?(2 votes)
- 1) What are the differences between energetic stability and thermodynamics?
2) What does it mean by kinetically stable ?
3) Is activation energy needed to overcome the kinetic energy barrier OR is activation energy the kinetic energy barrier itself?(1 vote)
In this video, I want to go ahead and talk about thermodynamics and kinetics. I remember that when I was first learning about these two things in the context of chemical reactions, I used to hear phrases like, oh, this chemical reaction-- let's say hypothetically A going to B-- is kinetically favorable, so it must proceed. But then the next day when I go to class, I'd hear another statement that said, oh, well this reaction A going to B is thermodynamically favorable, so it must proceed. So in this video, I want to go ahead and clarify these two statements for you and really understand what it means to thermodynamically favored and/or kinetically favored. So to get it started, I've already drawn two plots so that we can plot out the free energy change that occurs for the forward reaction-- we'll do that on the left side-- and the reverse reaction, which we'll do on the right side. So before we get into that, let's go ahead and label our axes. Our y-axis in both cases will be measuring free energy, which is in units of joules. Then our x-axis will be an abstract dimension called the reaction coordinate, which essentially allows us to monitor the progress of a reaction. So now I'm going to go ahead and say that the forward reaction we'll draw out in this teal color. And the reverse reaction we'll go ahead and draw out in this pink color. So let's go ahead and start with the forward reaction. So I'm just going to go ahead and say that the forward reaction has a negative delta G value. So remember that means it's spontaneous, and visually that means that our reactants start off at a higher free energy than our products. Now, this of course means that the reverse reaction, which will have the same magnitude of delta G, that is the free energy change, will be at the same numerical value. But of course, since the reaction is going the opposite direction, the sign of delta G will be now positive. And visually, we're saying essentially that our reactant, which in this case is B, starts our at a lower free energy than our product, which is A. Now thus far in drawing this free energy diagram, we've just been talking about thermodynamics. But it turns out that there is also a kinetic energy barrier for the conversion of reactants to products, regardless of whether the reaction is spontaneous or non-spontaneous. This kinetic barrier of energy is referred to as the free energy of activation, or simply activation energy. So I'm going to go ahead and put in parentheses E sub A, which we'll say stands for activation energy. And remember that delta G of course is talking about thermodynamics. With that said, let's go ahead and add this kinetic energy barrier to our diagrams. And we can do this by understanding that the activation energy is defined as the amount of energy that is required to form a high-energy intermediate during the course of the reactions. In other words, in our hypothetical reaction of A going to B, it proceeds through an intermediate, that is a high-energy chemical product that won't last very long, but is important in the conversion of A to B or vice versa. So in our diagrams here, we can go ahead and indicate that there is some intermediate, so in between our reactants and our products, that is at much higher energy than everything else. And we can go ahead and then connect the dots. So go ahead and essentially draw a line from reactants to products that includes our high-energy intermediate. And this ultimately allows us to see the presence of the activation energy as well as the change in free energy. So let's go ahead and actually label these things. So on the left side here, remember our change in free energy, which is looking at our reactants compared to our products-- we're ignoring this little bump. The change in free energy, which I'm going to indicate with the green line, extends between these two points. And the same on the right side, just again extending between the start and endpoints of our reactions. So I'll go ahead and indicate that this green line in both cases refers to our delta G values. And in a different color, let's say red, I'm going to go ahead and indicate the activation energy, which takes into account the change in energy between the high-energy intermediate and the reactants. So in the case on the left, that change is indicated here with red. And on the right side, the change between the intermediate and the reactant is a bit longer, so we'll go ahead and indicate that here. Now, activation energy is an important quantity to take into account, because in order for molecules to react, they must have enough energy to overcome this activation energy barrier. Essentially, in the case of a spontaneous reaction for example, I think of it like the energy one needs to get a ball to start rolling down a hill. We all know that gravity will make a ball roll down a hill, which is like a negative delta G value, it's telling us that the reaction is very thermodynamically favorable. But we need to sometimes give the ball push in order for the reaction to occur. And so that's kind of this little help that it needs to go over before it can actually proceed. For a non-spontaneous reaction, the idea is essentially the same. We still need to have some activation energy. But in addition, because it requires an input of energy, we can think about it as rolling ball up a hill instead of down a hill. Now in general, the idea is that the lower this free energy change, the faster a reaction will occur. And remember I'm saying faster, so I'm talking about kinetics. I'm talking about the rate of a reaction. So just to write that out, the activation energy, the smaller it is, the faster the reaction will proceed. Now in biochemistry in particular, it's really important to distinguish between these two terms of thermodynamics and kinetics, which we've drawn out in our diagrams as the change in delta G over the change in activation energy. Because many biochemical reactions in our body are kinetically unfavorable, that is to say they have a very high energy of activation even if they are thermodynamically favorable. This is why our bodies have enzymes, which essentially lower the activation energy of a reaction. So I went ahead and drew a dotted white line that's a little bit lower, so you can see that when an enzyme is present, the height of the barrier has decreased. And if it's decreased, the reaction will proceed faster. Now, there's one analogy that my chemistry professor used to tell us all the time that really helped me understand the interplay between kinetics and thermodynamics as they apply to whether or not a reaction will occur. So I'm going to go ahead and scroll down so we can briefly talk about this analogy, which is I think a fun way to think about all of this. So let's say you went to a dating website, because you were looking for your perfect match. And this dating website told you that your perfect match lived halfway across the globe. And this is such a perfect match, and they have all these algorithms. And if this match were a chemical equation, we would say that you had a very, very negative delta G value. That means you would be very, very spontaneous. But you live halfway around the world, so in terms of chemistry language, we might say that you are kinetically limited. That is to say, you don't have any way of actually travelling halfway across the world to meet your special someone. So we might say that you have a very high activation energy. So from this discussion, perhaps the biggest take-away is that neither kinetics nor thermodynamics solely determines whether a reaction will proceed. It's important to take both into account. Now one of the applications of this, in biochemistry especially, is remember that enzymes lower the activation energy of a reaction. So we can think of enzymes as light switches. They can regulate whether a reaction will proceed or won't proceed. But, of course, a light switch only works if the light bulb itself is working. And so the metaphor for the working light bulb is saying that a reaction has a negative delta G value. Or if light bulb is not working, we're saying that the delta G value is positive. That is, it requires a new battery or energy for the reaction to proceed.