I was having trouble with this problem. initial pressure for the compounds involved in the reaction displayed were determined to be P(CO(g)) = 0.5794 atm, P(H2O(g)) = 0.5662 atm, P(CO2(g)) = 0.7950 atm, P(H2(g)) = 0.2754 atm. Calculate the value of the equilibrium constant (Kp) at 1650 °C if the equilibrium pressure of CO2(g) was 0.6830 atm.

CO(g)+H2O(g) = CO2(g)+H2(g)

This looks straight forward to me. What is you trouble? What values do you have for the equilibrium concns?

To calculate the equilibrium constant (Kp) for the given reaction, you need to use the partial pressures of the compounds involved. Here's how you can solve it step by step:

1. Write the balanced equation for the reaction:
CO(g) + H2O(g) = CO2(g) + H2(g)

2. Set up the expression for Kp using the partial pressures:
Kp = (P(CO2) * P(H2)) / (P(CO) * P(H2O))

3. Before plugging in the values, convert the pressures from atm to bar (1 atm = 1.01325 bar):
P(CO) = 0.5794 atm * 1.01325 bar/atm ≈ 0.5868 bar
P(H2O) = 0.5662 atm * 1.01325 bar/atm ≈ 0.5736 bar
P(CO2) = 0.6830 atm * 1.01325 bar/atm ≈ 0.6929 bar
P(H2) = 0.2754 atm * 1.01325 bar/atm ≈ 0.2791 bar

4. Substitute the values into the Kp expression:
Kp = (0.6929 * 0.2791) / (0.5868 * 0.5736)
Kp ≈ 0.515

Therefore, the value of the equilibrium constant (Kp) at 1650 °C is approximately 0.515.