There are two flask, one filled with a)ice bath, and b)boiling water.
Ice bath: 1.00 L O^2(g) at STP
Boiling water: 1.00 O^2(g)
If we want the pressure to remain at 1.00bar when the O^2(g) is heated to 373K, what mass of O^2 must we release from the flask?
I know that I should use the general gas equation, but I don't know what quantities I should plug in!
I assume this is O2 at 0C and O2 at 100 C and it ISN'T dissolved in H2O.
Use PV = nRT. I would substitute P in atm (convert bar to atm), R = 0.08206 L*atm/mol*K and T = 273. Solve for n.
Do the same for the 373 system and solve for n.
Subtract to find the difference and convert mols to grams.
To find the mass of O^2 that must be released from the flask to maintain a pressure of 1.00 bar when heated to 373K, you can follow these steps:
Step 1: Determine the initial moles of O^2 in the flask at STP.
Since the flask is filled with 1.00 L of O^2 at STP, you can use the ideal gas law equation to find the moles:
PV = nRT, where P is the pressure, V is the volume, n is the number of moles, R is the ideal gas constant, and T is the temperature.
At STP, the pressure (P) is 1.00 bar, the volume (V) is 1.00 L, and the temperature (T) is 273.15K (standard temperature). The ideal gas constant (R) is 0.0821 L·atm/(mol·K) or 8.31 J/(mol·K).
Rearranging the ideal gas law equation to solve for moles (n), you get:
n = PV / RT
Substituting the values, you can calculate the initial moles of O^2 in the flask.
Step 2: Calculate the final moles of O^2 required to maintain a pressure of 1.00 bar at 373K.
Now, you need to find the number of moles of O^2 required for the final temperature of 373K. Use the same ideal gas law equation:
n = PV / RT
Here, the pressure (P) is still 1.00 bar, the volume (V) remains the same (1.00 L), the temperature (T) is now 373K, and the ideal gas constant (R) is the same.
Step 3: Find the difference in moles between the initial and final states.
Subtract the initial moles of O^2 (from Step 1) from the final moles of O^2 (from Step 2) to find the difference.
Step 4: Convert the difference in moles to mass.
To convert the difference in moles to mass, you need to use the molar mass of O^2, which is 32.00 g/mol. Multiply the difference in moles by the molar mass to get the mass of O^2 released from the flask.
By following these steps and plugging in the appropriate quantities into the ideal gas law equation, you can calculate the mass of O^2 that needs to be released from the flask.