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

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!!

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.

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 XN2O4 = PN2O4/total P
XNO2 found similarly.

Sure, let's go through each part step by step:

a) The balanced equation for the reaction is indeed N2O4 --> 2NO2.

b) To find the pressure if the gas in the flask at 298K is all N2O4, you can use the ideal gas law equation:

PV = nRT

We know the volume (V) is 5L, the temperature (T) is 298K, and since the gas in the flask is all N2O4, we want to find the number of moles (n) of N2O4. To find the number of moles, we can use the molar mass of N2O4, which is 92 g/mol.

n = mass/molar mass = 43.8 g/92 g/mol = 0.47608 mol

Now we can substitute the known values into the ideal gas law equation:

P(5L) = (0.47608mol)(0.0821 L.atm/mol.K)(298K)

P = (0.47608mol)(0.0821 L.atm/mol.K)(298K) / 5L
P ≈ 2.56 atm

So the pressure should be approximately 2.56 atm, not 2.33 atm.

c) To find the pressure if all the gas in the flask is converted into NO2, we can use the ideal gas law equation again. This time, we want to find the number of moles (n) of NO2. We can use the molar mass of NO2, which is 46 g/mol.

n = mass/molar mass = 43.8 g/46 g/mol ≈ 0.952 mol

Now we can substitute the values into the ideal gas law equation:

P(5L) = (0.952mol)(0.0821 L.atm/mol.K)(298K)

P = (0.952mol)(0.0821 L.atm/mol.K)(298K) / 5L
P ≈ 4.65 atm

So the pressure should be approximately 4.65 atm, which matches the given answer.

d) To find the mole fractions of N2O4 and NO2, we need to calculate the total number of moles in the flask at the stabilized pressure of 2.96 atm. We can use the ideal gas law equation again, this time solving for the total number of moles (n).

P(5L) = n(0.0821 L.atm/mol.K)(298K)

n = (2.96 atm)(5L) / (0.0821 L.atm/mol.K)(298K)
n ≈ 0.6 mol

Now we can find the mole fraction of N2O4 and NO2. The mole fraction (x) is given by:

x = n_component / n_total

Mole fraction of N2O4 = n_N2O4 / n_total = (0.47608 mol) / (0.6 mol) ≈ 0.793

Mole fraction of NO2 = n_NO2 / n_total = (0.952 mol) / (0.6 mol) ≈ 1.587

So the mole fractions of N2O4 and NO2 in the flask at the stabilized pressure of 2.96 atm are approximately 0.793 and 1.587, respectively.

I hope this helps clarify the steps! Let me know if you have any further questions.

To solve these questions, we need to apply the Ideal Gas Law, which states that the product of pressure (P) and volume (V) is directly proportional to the number of moles (n), multiplied by the gas constant (R) and temperature (T). The equation is given as:

PV = nRT

Where:
P = pressure
V = volume
n = number of moles
R = gas constant
T = temperature

Now let's solve each part of the question step by step:

a) The balanced equation is correctly given as N2O4 → 2NO2. This means that one molecule of N2O4 dissociates to form two molecules of NO2.

b) In this part of the question, we are assuming that all the gas in the flask is N2O4. We can calculate the pressure using the Ideal Gas Law:

PV = nRT

First, let's calculate the number of moles (n) of N2O4. The molar mass of N2O4 is 92 g/mol, so:

n = mass / molar mass
n = 43.8 g / 92 g/mol
n = 0.476 moles

Now plug in the values into the equation:

P * V = n * R * T

V is given as 5 L, R is the gas constant (0.0821 L·atm / mol·K), and T is given as 298 K. Solving for P:

P = (n * R * T) / V
P = (0.476 moles * 0.0821 L·atm / mol·K * 298 K) / 5 L
P ≈ 2.67 atm

c) In this part of the question, we assume that all the N2O4 has converted into NO2. Since we already know the number of moles of N2O4 (0.476 moles), we can use stoichiometry to calculate the number of moles of NO2 formed. From the balanced equation, 1 mole of N2O4 gives 2 moles of NO2. Therefore:

moles of NO2 = moles of N2O4 * 2
moles of NO2 = 0.476 moles * 2
moles of NO2 = 0.952 moles

Using the Ideal Gas Law, we can calculate the pressure:

P = (n * R * T) / V
P = (0.952 moles * 0.0821 L·atm / mol·K * 298 K) / 5 L
P ≈ 4.65 atm

d) To find the mole fractions of N2O4 and NO2, we need to know the total number of moles of gas present when the pressure stabilizes at 2.96 atm.

First, let's calculate the total number of moles using the Ideal Gas Law:

P * V = n * R * T

Rearrange to solve for n:

n = (P * V) / (R * T)
n = (2.96 atm * 5 L) / (0.0821 L·atm / mol·K * 298 K)
n ≈ 71.29 moles

Now, we can calculate the mole fraction of N2O4:

mole fraction of N2O4 = moles of N2O4 / total moles of gas
mole fraction of N2O4 = 0.476 moles / 71.29 moles
mole fraction of N2O4 ≈ 0.0067

Similarly, the mole fraction of NO2 can be calculated:

mole fraction of NO2 = moles of NO2 / total moles of gas
mole fraction of NO2 = 0.952 moles / 71.29 moles
mole fraction of NO2 ≈ 0.0133

To summarize:
b) The pressure when all gas in the flask is N2O4 is approximately 2.67 atm.
c) The pressure when all gas in the flask is NO2 is approximately 4.65 atm.
d) The mole fractions of N2O4 and NO2, once the pressure stabilizes at 2.96 atm, are approximately 0.0067 and 0.0133, respectively.