A flask has 1.00 L of O2 (g) at STP (the amount of gas and volume are held constant, P=1 bar) and the second has 1.00 L of O2 at 100 C. What mass of O2 must release if we get the pressure to remain at 1.00 bar when it's heated to 100 C?
Can someone please explain how to do this? What i tried to do was calculate the n of the first flask at STP. I got .00403mol. I then calculated the n of the second flask using n(initial)T(initial)/T(final)
I'm not sure where to go from here. I'm also not sure whether I calculated it correctly. Any help would be appreciated, thank you
I may have missed something but I don't know what the second flask has to do with the problem. I would use PV = nRT and solve for mols in the first flask. Then use PV = nRT but this time with 100 C, then subtract the two n values.
To solve this problem, you need to use the ideal gas law equation:
PV = nRT
Where:
- P is the pressure of the gas (in this case, 1.00 bar)
- V is the volume of the gas (in this case, 1.00 L)
- n is the number of moles of gas
- R is the ideal gas constant (0.0821 L·atm/(mol·K) or 8.314 J/(mol·K))
- T is the temperature of the gas (in this case, given in Celsius)
Let's break down the steps:
Step 1: Calculate the number of moles of O2 in the first flask at STP:
Given that the amount of gas and volume are held constant, and the flask is at STP, we can use the ideal gas law to find the number of moles (n) of O2:
PV = nRT
n = PV / RT
Substituting the given values:
P = 1.00 bar
V = 1.00 L
T = 273.15 K (STP temperature, 0°C or 32°F)
Now, calculate n using the ideal gas law equation. Make sure to convert the pressure to atmospheres (1 bar = 0.987 atm):
n = (1.00 atm) * (1.00 L) / (0.0821 L·atm/(mol·K) * 273.15 K)
This gives you the number of moles of O2 in the first flask at STP.
Step 2: Calculate the number of moles of O2 in the second flask at 100°C:
This step requires a slightly different approach. You need to apply the temperature change to find the number of moles of O2 in the second flask, while keeping the pressure constant at 1.00 bar.
Use the equation:
n(final) = n(initial) * (T(initial) / T(final))
Given:
n(initial) = the number of moles of O2 in the first flask at STP (obtained in Step 1)
T(initial) = the temperature of the first flask (STP temperature, 273.15 K)
T(final) = the final temperature of the second flask (100°C = 373.15 K)
Now, substitute the values and calculate n(final).
Step 3: Calculate the mass of O2 released:
To find the mass of O2 released, you will need to subtract the number of moles in the second flask from the number of moles in the first flask. Then, calculate the mass using the molar mass of O2 (M = 32.0 g/mol).
Mass = (n(initial) - n(final)) * molar mass of O2
Now, you can substitute the values and calculate the mass of O2 released.
Note: Make sure to convert between Celsius and Kelvin when applying temperature values in the equations.
I hope this explanation helps!