Having some confusion with this question....I have the answers, just not able to reach them.

Question

A flask of volume 5L is evacuated and 43.8 g of solid N2O4 is introduced at 77K. The sample is then warmed to 298K during which the N2O4 vaporizes and some dissociates to form brown NO2 gas. The pressure slowly increases until it stabilizes at 2.96 atm.

a) balanced equation (Which I know is N2O4 --> 2NO2)

b) if gas in the flask at 298K is all N2O4 what would the pressure be?
--answer is 2.33, but I keep getting about 2.6

c) if all gas in flask converted into NO2, what would pressure be? Answer is 4.65 atm which I did find...however I'm not sure if my method was just a fluke since I tried the same/similar thinking with part b, and can't get the right answer

d) what are the mole fractions of N2O4 and NO2 once the pressure stabilizes at 2.96 atm?

for c I found the molar mass of NO2= 46 g/mol then divided that into 43.78 g to get .952 moles. I then used the PV=nRT for P to find 4.65 atm.

As for the others mentioned...not sure what I should be doing...I'm pretty confused at this point. Thanks!!

Answer 1

Do you have a Kc or Kp?
b)P = nRT/V = 0.476*0.08206*298/5 = 2.328 which rounds to 2.33 atm.

c)Actually, it is just 2*PN2O4 because mole NO2 = 2*moles N2O4. 2*2.33 = 4.66 atm. or 2*0.470*0.08206*298/5 = 4.656 which rounds to 4.66 atm.

Answer 2

d)
I'm thinking that if we call the partial pressure N2O4 = x, then partial pressure NO2 is 2x and we know
x + 2x = 2.96. Solve for x.
Then X_N2O4 = P_N2O4/total P
X_NO2 found similarly.