Question

part a

Pure NO2(g) decomposes at 1000K into NO(g) and O2(g) according to the reaction: 2NO2(g) ⇌ 2NO(g) + O2(g). At 1000K, the equilibrium constant is Kp = 158. If analysis shows the equilibrium partial pressure of O2 to be 0.26 atm, what is the equilibrium pressure of NO? Assume units of atm.

part b

For the reaction: N2(g) + O2(g) ⇌ 2NO(g) If the equilibrium partial pressures for N2, O2 and NO are 0.15 atm, 0.33 atm and 0.05 atm, respectively, at a particular temperature, what is Kp?

part c

Pure NO2(g) decomposes at 1000K into NO(g) and O2(g) according to the reaction: 2NO2(g) ⇌ 2NO(g) + O2(g)

At 1000K, the equilibrium constant is Kp = 158. If analysis shows the equilibrium partial pressure of O2 to be 0.26 atm, what is the equilibrium pressure of NO2?

part d

For the following reaction in a closed container at 351°C: CaCO3(s) ⇌ CaO(s) + CO2(g). The partial pressure of CO2 is 0.12 atm. Assume this reaction is the only source of CO2. What is Kc for this reaction? Assume units of mol/L

Answer 1

Part a )The reaction is 2NO₂(g) ⇌ 2NO(g) + O₂(g)

When O₂ is 0.26 atm produced, twice as many moles of NO are produced = 0.26 x 2 = 0.52 atm NO

Part b) The given reaction is---  N₂(g) + O₂(g) ⇌ 2NO(g)

Now, K_p = [NO]² / [N₂] [O₂]

=> K_p = (0.05)² / 0.15 x 0.33

=> K_p = 0.0025/0.0495

=> K_p = 5.1 x 10^-2 or 0.051

Part c) The reaction is 2NO₂(g) ⇌ 2NO(g) + O₂(g)

When O₂ is 0.26 atm produced, twice as many moles of NO are produced = 0.26 x 2 = 0.52 atm NO

Now, K_p = [NO]² [O₂] / [NO₂]²

=> 158 = (0.52)² x 0.26 / [NO₂]²

=> [NO₂]² = (0.52)² x 0.26 / 158

=> [NO₂]² = 0.070304/158

=>  [NO₂]² = 0.00044496

=> [NO₂] = 0.00044496

=> [NO₂] = 0.021 atm.