A mixture of CO(g) and O₂(g) in a 1.0-L container at 1.0×10³ K has a total pressure of 2.2 atm. After some time, the total pressure falls to 1.9 atm as the result of the formation of CO₂. Determine the mass (in grams) of CO₂ that forms.

Verified step by step guidance

1. Identify the initial and final conditions of the system: The initial total pressure is 2.2 atm, and the final total pressure is 1.9 atm.
2. Determine the change in pressure due to the formation of CO2: Subtract the final pressure from the initial pressure to find the pressure change.
3. Use the ideal gas law to relate the change in pressure to the change in moles of gas: \( \Delta P = \Delta n \frac{RT}{V} \), where \( R \) is the ideal gas constant, \( T \) is the temperature, and \( V \) is the volume.
4. Calculate the change in moles of gas (\( \Delta n \)) using the pressure change and the ideal gas law equation.
5. Convert the change in moles of CO2 to mass using the molar mass of CO2 (44.01 g/mol).

Key Concepts

Ideal Gas Law

The Ideal Gas Law relates the pressure, volume, temperature, and number of moles of a gas through the equation PV = nRT. In this scenario, understanding how changes in pressure and temperature affect the number of moles of gas is crucial for determining the amount of CO2 produced from the reaction.

Stoichiometry

Stoichiometry involves the calculation of reactants and products in chemical reactions based on balanced equations. In this case, knowing the stoichiometric relationships between CO, O2, and CO2 will help determine how much CO2 is formed as the reaction proceeds and the pressure changes.

Partial Pressure

Partial pressure is the pressure exerted by a single component of a gas mixture. According to Dalton's Law, the total pressure of a gas mixture is the sum of the partial pressures of its components. Understanding how the total pressure changes as CO2 forms will allow for the calculation of the amount of CO2 produced.