To determine the limiting reactant, we need to calculate the amount of CO2 that can be produced from each reactant and see which reactant produces less CO2. Then, we can use the amount of CO2 produced by the limiting reactant to calculate the total pressure.
The balanced equation tells us that 2 moles of CO react with 1 mole of O2 to produce 2 moles of CO2. Therefore, we can calculate the amount of CO2 that can be produced from the given amounts of CO and O2:
CO: 1.5 moles CO × (2 moles CO2 / 2 moles CO) = 1.5 moles CO2
O2: 2.0 moles O2 × (2 moles CO2 / 1 mole O2) = 4.0 moles CO2
Since the amount of CO2 produced from CO (1.5 moles) is less than the amount of CO2 produced from O2 (4.0 moles), CO is the limiting reactant.
To calculate the partial pressures of CO2 and O2, we need to use the ideal gas law:
PV = nRT
where P is the pressure, V is the volume, n is the number of moles, R is the gas constant, and T is the temperature.
Assuming ideal gas behavior, we can use the ideal gas law to calculate the partial pressure of CO2 produced from the reaction:
n(CO2) = 2 × n(CO) = 2 × 1.5 = 3.0 moles
P(CO2) = n(CO2) × RT/V = 3.0 × 0.0821 × 273/10.0 = 6.31 atm
To calculate the partial pressure of O2, we need to subtract the amount of O2 consumed by the reaction from the initial amount:
n(O2) = 2.0 - 2 × n(CO) = 2.0 - 2 × 1.5 = -1.0 moles (negative because O2 is in excess)
Since the moles of O2 is negative, it means there is no O2 left, and therefore its partial pressure is zero.
The total pressure in the flask is the sum of the partial pressures of CO2 and O2:
P(total) = P(CO2) + P(O2) = 6.31 + 0 = 6.31 atm
Therefore, the total pressure in the flask at 0.0 °C is 6.31 atm.