A 10 liter flask at 298 K contains a gaseous mixture of O2 and CO2 at 1 atmosphere. Which statement is true for the partial pressures of O2 and CO2 if 0.2 mole of O2 is present in the flask? (Given the universal gas constant R = 0.082 L∙atm/K∙mol)
You didn't give the options for an answer but here is how you determine the parital pressure of O2. It can also be solve another way by going through mole fractions.
Use PV = nRT. Plug in V, R, and T (in kelvin remember) and use 0.2 mole for n. Solve for P
To determine the partial pressures of O2 and CO2 in the flask, we can use the ideal gas law equation:
PV = nRT
Where:
- P represents the pressure (in atmospheres)
- V represents the volume (in liters)
- n represents the number of moles
- R represents the universal gas constant
- T represents the temperature (in Kelvin)
Let's use this equation to find the partial pressure of O2:
1) Convert the temperature from degrees Celsius to Kelvin:
298 K = 298 K
2) Calculate the total number of moles of gas in the flask:
Given that 0.2 moles of O2 is present, the total number of moles will be 0.2 moles (since there is no information about the number of moles of CO2).
3) Plug in the values into the ideal gas law equation for O2:
P(O2) * V = n(O2) * R * T
Since the volume (V) is given as 10 liters, and the temperature (T) is given as 298 K, we can rearrange the equation:
P(O2) = (n(O2) * R * T) / V
Substituting the values we have:
P(O2) = (0.2 moles * 0.082 L∙atm/K∙mol * 298 K) / 10 L
Simplifying the equation:
P(O2) = 0.00492 atm
So, the partial pressure of O2 in the flask is 0.00492 atm.
Since we don't have information about the number of moles of CO2 present, we cannot determine its partial pressure.