For the following reaction, the partial pressures are listed in the table:
Substance Equlibrium
H2S 5.4
I2 1.2
HI 0.43
S 7
If the Kp of the reaction is 0.134, which direction would the reaction need to go to establish equilibrium?
H2S(g) + I2(s) <--> 2HI(g) + S(s)
The correct answer is: "Since Q<Kp, the reaction must proceed to the right."
Now there are two roberts. Palm to face
So is the answer that Qp and Kp are equal?
at equilibrium, they are equal
To determine the direction in which the reaction needs to go to establish equilibrium, we need to compare the given partial pressures with the equilibrium constant (Kp) of the reaction.
The equilibrium constant (Kp) expression for the reaction is:
Kp = (P(HI)^2 * P(S)) / (P(H2S) * P(I2))
Given partial pressures:
P(H2S) = 5.4
P(I2) = 1.2
P(HI) = 0.43
P(S) = 7
Substituting these values into the Kp expression:
Kp = (0.43^2 * 7) / (5.4 * 1.2)
Kp = 2.7647
Since the given Kp value (0.134) is smaller than the calculated Kp value (2.7647), it means that the reaction has more products than is predicted by the given partial pressures.
To establish equilibrium, the reaction needs to favor the formation of reactants and therefore shift to the left-hand side (towards the reactants). This means that the reaction needs to go in the reverse direction:
2HI(g) + S(s) <--> H2S(g) + I2(s)
By shifting to the left side, the concentrations of H2S and I2 will increase, while the concentrations of HI and S will decrease, ultimately resulting in a new equilibrium position where the ratio of the concentrations will be consistent with the given Kp value.