This question explores Le Chatelier’s principle applied to the equilibrium between hydrogen, nitrogen, and ammonia:
N2(g) + 3 H2(g) ⇌ 2 NH3(g) Kp = 5.8 × 106 atm−2 at 25 °C
a) Write the expression for Kp for this reaction.
b) Predict the effect on the equilibrium (at 25 °C) of an increase in each partial pressure:
c) Suppose that 0.245 mol N2, 0.00145 mol H2, and 0.162 mol NH3 occupy a 10.0-L volume; the partial pressures are PN2= 0.600 atm, PH2= 0.00356 atm, and PNH3= 0.396 atm and the total pressure is 1.00 atm. Is the system at equilibrium? If not, in which direction would the reaction shift to reach equilibrium? Show a calculation to support your answer.
d) Suppose the total pressure is kept constant and that 0.010 mol N2 is added to the system described in part c so that the total amount of N2 is 0.255 mol and the total amount of all gases is 0.418 mol. If the total pressure stays constant, what must happen to the volume? Why?
e) Based on the ideal gas law it is possible to calculate the new volume of the mixture of gases and to calculate the partial pressures of the constituent gases. The partial pressures are PN2 = 0.610 atm, PH2 = 0.00347 atm, and PNH3 = 0.388 atm. Is the system at equilibrium? If not, in which direction does the reaction shift to reach equilibrium? Show a calculation to support your answer.
f) Does your answer in part (e) agree with your first answer in part b? Explain why or why not.
g) Have you discovered an exception to Le Chatelier’s principle? Explain why or why not.
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