00:01
This question is dealing with equilibrium.
00:04
And in this case, we have this reaction.
00:09
I'm not told all of the information that we need right within those questions.
00:14
I had to look up some of the information, but here's our balanced reaction.
00:20
And we're told that at 2000 kelvin is when the reaction is going to be taking place.
00:28
The kp or equilibrium pressure constant is 0 .061 .1.
00:32
I had to look up this piece of information, but here it is.
00:36
Now, knowing that and that the initial partial pressures for each of the two reactants are these values, let's figure out the partial pressures of the products at equilibrium.
00:46
So we set up an ice table, knowing that the initial pressures i've plugged in for the carbon monoxide and water reactants, and before anything takes place, of course, the concentrations of the two products are going to be zero.
01:02
So over time, there's a one -to -one ratio between each of the products and reactants here.
01:09
So over time, as x amount of the reactant is lost, that same amount of product will be produced, so that at equilibrium, these are my concentrations, 1344 minus x, 1766 minus x for the two reactants, and x the partial pressure or concentration for each of the two products.
01:30
Now, if i look at the equilibrium constant for pressure, the partial pressure of the co2 and hydrogen products over the partial pressure of the two reactants is going to equal that kp value of 0 .0611.
01:46
So let's plug in what we know now about those partial pressures.
01:51
So at equilibrium, partial pressure of the carbon dioxide product is x, and so is the hydrogen product.
02:01
So we'll plug that value in, and the equilibrium concentration of the carbon monoxide is 1344 minus x, 1766 minus x for the water product.
02:16
And again, this value is going to be equal to the 0 .0611 kp value that we were given...