Given the water shift reaction:
CO(g)+H2O(g)⇌CO2(g)+H2(g) such that Kp = 4.0 at 500 °C
and the initial concentrations of 0.100 M CO and 0.100 M H2O, we should be able to find the partial pressures of each component at equilibrium
I wanted to check if my logic was correct in attempting this problem: I assumed the reaction took place in a 1 L container, yielding 0.100 mols of CO and H2O. I then used the ideal gas law to determine the partial pressures of CO and H2O to both be about 6.34 atm.
Then I can use the ICE chart to determine the equillibrium pressures to be 6.34 - x for CO and H2O and just x for the products.
4.0 = (x^2)/(6.34-x)^2
by the quadratic equation, x = 4.22 or x ≈ 12
The x=4.2 solution is the only one that makes sense because x = 12 would yield a negative answer for pressure which is impossible in this context
Thus, the partial pressures are
P(CO) = P(H2O) = 6.34 - x
= 2.1 atm
P(H2) = P(CO2) = x
= 4.2 atm
Is this process of doing the problem correctly? My professor won't release the answer key until after the homework set is due and I couldn't find this in any textbook