Consider the reaction 2 NO2 (g) ⇌ 2 NO (g) + O2 (g). If the initial partial pressure of NO2 (g) is 3.0 bar, and x is the equilibrium pressure of O2 (g), what is the correct equilibrium relation?

..........2NO2(g) ⇌ 2NO(g) + O2(g)

I..........3..........0.......0
C.........-2x........ 2x......x
E.........3-2x....... 2x......x

OK?

por que -2x si lo que se le quita a No2 es -2x^2

To determine the equilibrium relation for the given reaction, we need to use the balanced equation and the concept of equilibrium expressions.

The balanced equation for the reaction is: 2 NO2 (g) ⇌ 2 NO (g) + O2 (g)

The equilibrium expression for the reaction can be written as follows:

Kp = (P(NO))^2 * P(O2)

Where Kp is the equilibrium constant, P(NO) is the partial pressure of NO, and P(O2) is the partial pressure of O2.

In the given reaction, the initial partial pressure of NO2 (g) is 3.0 bar. Let's assume the equilibrium pressure of NO (g) is y bar, and the equilibrium pressure of O2 (g) is x bar.

So, we can write:

P(NO2) = 3.0 bar

P(NO) = 2y bar

P(O2) = x bar

Substituting these values into the equilibrium expression, we get:

Kp = (2y)^2 * x

Now, to determine the correct equilibrium relation, we need to compare the value of Kp with the actual equilibrium expression.

Unfortunately, the value of Kp is not provided in the given question. Thus, we cannot determine the exact relationship between the pressures of NO and O2 without the value of Kp.

However, we can make a general statement. If Kp is a large value, it indicates that the forward reaction is favored, which means the concentrations of the products (NO and O2) are higher compared to the reactant (NO2) concentration at equilibrium. On the other hand, if Kp is a small value, it indicates that the reverse reaction is favored, which means the concentration of the reactant (NO2) is higher compared to the concentrations of the products (NO and O2) at equilibrium.

So, without the specific value of Kp, we cannot determine the equilibrium relation accurately.