In this module, we are going to look at chemical equilibrium, and we'll start by looking at some definitions. So when we talk about equilibrium, what we do is we look at a system, which system is underlined here. It means something really specific with how we frame things.
We won't have to worry about that too much until we get into a later module. Really, our system is probably a reaction. So we're going to look at this system, and it will be in one of two states.
It's either going to be changing or not changing. And if we wait, if something is changing and we wait an appropriate amount of time, it will get to a state where it's not changing. And when it is not changing, this is what we consider to be equilibrium.
Now we need some equilibrium definitions. So this is when macroscopic qualities don't change. And we're talking about color, pressure, concentration, things where if one molecule of product turned back into a reactant at the same time that one molecule of reactant or whatever a reaction looks like turns into product, there's still some change happening. But overall, the concentration wouldn't change, the color wouldn't change, the pressure wouldn't change. Another definition, this is related to kinetics, which we will look at much later, but it's an important definition for equilibrium.
And that's when the rate of the forward reaction equals the rate of the reverse reaction. So it doesn't mean that everything's perfectly balanced, but the rate at which we're making product is the same as the rate at which our product is turning back into reactants, which is why we overall don't see a change. So...
I mean, let's put some data, put some graphs to this. We are going to look at a reaction achieving equilibrium. And the reaction that we're going to look at is 2NO2 gas in equilibrium with N2O4 gas.
Now, the reaction arrow here is special. This means equilibrium, where you have a forward arrow and a reverse arrow. It's telling us that the reaction is reversible.
And when we talk about the components of this reaction, just like way back when we talked about reactions for the first time, and I said, the stuff on the left will always be called the reactants. This is where it matters. So if it's on the left, if it's a reactant, if it's on the right, it's a product.
And if N2O4 is turning into 2NO2, because of how we have framed, we've written this reaction down in a specific way, that is a product turning into a reactant. It's the reverse of the reaction that we're looking at. And that's extremely important. Once we put our reaction on here in the order that it's in, with specific things on the left and specific things on the right, that's set whether they're reactants or products.
And this matters because the values that we're going to use to talk about equilibrium, equilibrium. change when the reaction changes. So we're going to keep our reference stable.
So experiment one that we're going to do with this reaction is we're going to start with one atmosphere of NO2 at 298 Kelvin, which means that's the only thing that we're starting with. So we're going to start with all reactants, no products. And what this would look like on my graph, I'm going to graph pressure versus time. So as my reaction goes, I start with one atmosphere of NO2.
I start with zero atmospheres of N2O4. And as my reaction proceeds, my pressure of NO2 goes down because it gets used up as I make product. My pressure of my product goes up from zero. And eventually, they reach a point at which...
their concentrations or their pressures don't change. So I've got this like dashed yellow line with a star. This is the point at which my pressure of N2O4 and my pressure of NO2 are constant. Now we could do another experiment. And in our other experiment, we're going to start with 0.5 atmospheres of N2O4 at 298K.
So now I'm starting with 0.5 atmospheres of product and zero. Reactant. But it doesn't matter because my reaction will proceed.
It just happens to proceed to the left this time because I only have product. So I've got to make some reactant. So my product pressure goes down as the reaction proceeds. My reactant pressure goes up from zero.
And there's a point at which those pressures stop changing. So I have that also marked on this graph with... dashed line and a yellow highlight.
So this is the point at which the equilibrium pressure is reached. And I have interchanged concentration a couple of times with pressure in describing this. Well, if you look at the ideal gas law, concentration is really just number of moles in a volume.
So we can relate concentration to pressure directly. Now we call this dynamic equilibrium. Because the reaction is still happening, we still have a little bit going forward, a little bit going backward, but because they proceed at the same rate, the overall pressures and overall concentrations stay the same. Now, some reactions achieve equilibrium with more product, and some achieve it with more reactant.
So it's not going to be an equal amount of each. The conditions and the characteristics of a specific Reaction will determine where that end point lies. So what we need is a concise way to describe the equilibrium state. And we use KEQ to do this. KEQ is the equilibrium constant.
You will see we'll kind of drop the EQ off of here and just write K in a lot of cases. There will be some helpful subscripts that come up. And sometimes we'll add EQ back in here just because we use the letter K for so many things.
But this capital K, capital K EQ is the equilibrium constant. So the way that this is expressed, if I have a generic reaction, which is A and B in equilibrium with C and D, and they have these coefficients, the equilibrium constant expression is A. The concentrations of the products raised to their coefficients divided by the concentrations of the reactants raised to their coefficients. It's always products over reactants. So in my generic reaction, this means it's the concentration of C to the little c power, concentration of D to the little d power, divided by concentration of A to the little a power times concentration of B to the little b power.
Okay, so a reminder, like that brackets mean concentration. We're doing, if it's concentration of C, right, it's moles of C divided by volume in liters. Moles per liter is molarity. And we will do some of these problems with pressure. Let's do one with concentration.
So that's how we're starting to define this. And I have a combustion reaction for my example. So... If my reaction is C3H8 plus 502 in equilibrium with 3CO2 and 4H2O, my equilibrium constant expression, which if I ask you for the equilibrium constant expression, this is what I'm asking for. So it's the concentration of the product raised to the power of the coefficient.
So concentration of CO2 to the third power times concentration of water to the fourth power. Divided by the concentration of C3H8 times the concentration of O2 to the fifth power. All right, so there will be a lot of practice with these. Make sure that you are comfortable writing equilibrium constant expressions for reactions. So what does KEQ tell us about a reaction?
So for an example reaction, let's look at H2 reacting with Br2. And that's in equilibrium with 2 HBr. This has a Keq of 1 times 10 to the 19th at 298 K. So 1 times 10 to the 19th, that's a very large number. If I write my equilibrium constant expression, which you should always do to fit things into this model that we're looking at, it's the concentration of HBr squared.
over the concentration of H2 times the concentration of Br2. I mean, maybe this doesn't help you understand what's going on with the value at this point, but... It should, right?
I have a very big number. If I think about what that means in relation to the division that's going on here of products over reactants, a very large KEQ means that at equilibrium, I have more product than I do reactant, right? To get a big number, the number in the numerator needs to be bigger than the number in the denominator when I divide. So we would describe this reaction as favoring the products.
Right, this is what a large value for Keq means. So another example, N2 plus O2 in equilibrium with 2NO. This is an equilibrium constant of 4.1 times 10 to the minus 31 at 298K. So a very small Keq.
My equilibrium constant expression is concentration of NO squared over concentration of N2 times concentration of O2. So a very small number means I have a larger denominator, the numerator, so I have more reactant at equilibrium. And we might say that this reaction favors the reactants.
Overall, what you will see is that if the equilibrium constant is greater than one, it favors the products. If it's less than one, it favors the reactants. Now to different degrees, right?
Like An equilibrium constant of 0.9 is very different from an equilibrium constant of 4.1 times 10 to the minus 31. But both of those are less than 1. So the balance would lie with the reactants at equilibrium. If your KEQ is exactly equal to 1, then, I mean, neither is favored. They're just equal in their balance between products and reactants.