Explore how we describe a chemical reaction using an energy diagram. So every time we have a diagram, we need to have an axis. The y-axis is going to have a unit of energy, and it's going to be in the form of the Gibbs free energy.
And on the x-axis, we're going to have a reaction progress. So let's first draw... a spontaneous chemical reaction. And remember, a spontaneous chemical reaction means that the delta G, the change in the Gibbs free energy, is going to be a negative value.
So in order for the delta G to be negative, the reactant needs to be in higher energy compared to the product. So it goes down in energy, right? So the reactant is higher energy compared to the product.
So when the chemical reaction happened, the delta G is negative. And what we extract as a form of usable work is the difference in this energy. So in this case, whenever you see a reactant that's higher than the product, the delta G is negative, and it's going to be a spontaneous process. So where does the activation energy that we just discussed earlier comes into play?
Well, an activation energy means that the two reactants or three or whatever it is need to collide faster than normal in order for this chemical reaction to happen such that the electrons inside of them would rearrange themselves. So we describe that by essentially adding a hill to climb, right? So that's a hill to climb. And the activation energy will start from the reactant side here.
So the activation energy of going from the reactant to the product side would be this one. Oops, I just erased more things than I wanted there. This is the activation energy. the forward reaction okay and this activation energy is going to be a positive value because you know we need to climb so we need to to expect it's it's an energy we need to overcome this energy so the energy is higher so the activation energy of the forward direction is going to go from here all the way to up there and then what I meant by the forward direction is that we're going to go from reactant to the product side, we have to overcome a barrier of this height, a hill of this height. So that's going to be the activation energy, and again for all spontaneous reactions, the reactant is going to be higher in energy compared to the product.
So the activation energy is going to, excuse me, the delta G is going to be this particular magnet, right? So the delta G is negative there. So that's going to be a picture of a spontaneous chemical reaction.
Now how about if we have a non-spontaneous chemical reaction? Well, the picture is going to be similar, but a non-spontaneous reaction means a delta G is positive. So this is Gibbs free energy, this is reaction. And again, for a non-spontaneous process, the delta G is positive. Non-spontaneous.
So what does it look like? Well, the reactant is going to be lower in energy now compared to the product. And the... delta G would be the difference between this energy from top to bottom. So the delta G would go from here all the way up here, right?
So that's the same level. So that's going to be our delta G, and this is going to be a positive value. So where's the picture of what the activation energy comes in?
Well, for all chemical reaction, you need to overcome a barrier one way or the other for a chemical reaction to happen. So let's say it looks something like that, yeah? Now, where's the activation energy? Well, the activation energy is going to start from the reactant side.
to the product sides and this is going to be the activation energy of the forward reaction activation energy energy of the forward reaction or let me just say the activation okay so we'll just same thing So that's the activation energy and the value is going to be positive as well because we have to climb up the hill of this height to go from the reactant to the top of the hill. All right. So that's pretty much it in terms of how we describe a chemical reaction using an energy diagram. Now, there's a couple of things that I want to mention. When we have a chemical reaction, that is, when your chemical reaction happens, the reactant and product is going to move.
forward and backward that is to say think of it as Each of your molecules that's inside your reaction, let's say we have a spontaneous reaction with a low enough energy barrier, let's say. So this is a reactant, this is a product, right? So what happens is that the chemical reaction, you know, atoms are moving around, so they have an energy. And sometimes you have some atom that's going to be able to move, overcome the barrier, and then go to the product side. But sometimes if you kind of assign the chemical reaction or whatever condition that you have, there's some small probability, but not completely eliminated, there's some possibility for a chemical reaction to go from the product side back to the reactant side.
Now, of course, for a chemical reaction that is spontaneous like this, sometimes we have a reactant and product, right? Sometimes if you have an equilibrium state, you kind of draw the arrow equally like this. But in reality, what we have if we have a spontaneous progress is that the forward direction going from reactant to product Is going to be faster than the reaction that's going backward, right? So that's kind of what we have now what happens is of course over time you have Equilibrium that is suppose that you have like 10 of these things right initially you have all of them in the reactant side So initially all the chemical reaction is just going to move forward But after a while what you'll find is that you're going to create enough product that you have a lot more product even though the small probably there's only small probability for each product to climb in you have a lot of them so what happens is that at equilibrium we draw it as equal arrow here because the reaction and product have now reached an equilibrium that is the amount of stuff going downward versus the amount of stuff going upward is going to be the same so that's going to be the picture at equilibrium now the last thing i want to mention about this whole thing is what happens if you have a reaction with a catalyst, because we discussed about catalysts on the previous slide, but we haven't actually talked about it at all. Now, what is a catalyst?
Well, a catalyst is a substance that speeds up the rate of chemical reaction, but it's not consumed in the reaction itself. So there's going to be an important aspect about a catalyst. It is not consumed. It's going to stay around.
Now, how does it change this kind of speeds up the chemical reaction? Well, typically, it substantially lowered the barrier along the reaction coordinate such that the activation energy is a lot smaller because you have a different pathway. So what does it look like?
Well let's say we have an exaggerated chemical reaction here. So this gives your energy reaction pathway. So we have a reactant and we have a product and let's say we have normally you have a huge activation energy like that right. So that's a huge activation energy.
Now this is the normal path and here's kind of the reaction with a catalyst you're going to add additional substance for organic chemistry this could be a platinum this could be a palladium this could be gold irons or a lot of different heavy metals can act as a catalyst and what happens is that the reactor is now going to use a different path so you're not going to go through the normal path and instead you're going to choose by adding this catalyst you're going to create a better pathway where the activation energy is left less. Let's say the activation energy is this, right? So that's kind of what we have. So this is pathway with catalyst. So the point is that the activation energy is going to be a lot less.
Now, of course, it's not always that clean. Sometimes your catalyst, if you have an enzyme, have a reaction pathway that actually looks like that. That's certainly possible. But in any sense, in essence, the important part is that the activation energy is less. So that's going to be the important part, right?
The pathway, path with catalyst and activation energy is lower. Okay, so that's going to be the important part about the catalyst. It lowers the activation energy by changing a different path, but it is not consuming the chemical reaction. So that's going to be an important part. And here's another important part.
The delta G of the chemical reaction remains the same. Delta G is the same whether you're using a normal path, uncatalyzed, or a catalyzed reaction. The delta G is the same. All right, so in the next video, I'm going to show you a combustion of methane gas on the normal path as well as the catalyzed path. And I want to try to see whether you can see where this mechanism difference between a catalyst and non-catalyzed reaction look like.