When 0.670 g argon is added to a 500 cm³ container with a sample of oxygen gas, the total pressure of the gases is found to be 1.52 atm at a temperature of 340 K. What is the mass of the oxygen gas in the bulb?

• A. 0.266 g
• B. 0.335 g
• C. 0.621 g
• D. 0.715 g
• E. 1.72 g

Answer: (B)

Verified step by step guidance

1. First, use the ideal gas law to find the partial pressure of argon. The ideal gas law is given by \( PV = nRT \), where \( P \) is pressure, \( V \) is volume, \( n \) is the number of moles, \( R \) is the ideal gas constant, and \( T \) is temperature.
2. Calculate the number of moles of argon using its mass and molar mass. The molar mass of argon is approximately 39.95 g/mol. Use the formula \( n = \frac{\text{mass}}{\text{molar mass}} \) to find \( n \).
3. Substitute the values for argon into the ideal gas law to find its partial pressure. Rearrange the ideal gas law to solve for \( P \): \( P = \frac{nRT}{V} \). Use \( R = 0.0821 \text{ L atm/mol K} \) and convert the volume from cm³ to L.
4. Subtract the partial pressure of argon from the total pressure to find the partial pressure of oxygen. The total pressure is given as 1.52 atm, so \( P_{\text{oxygen}} = P_{\text{total}} - P_{\text{argon}} \).
5. Use the ideal gas law again to find the number of moles of oxygen using its partial pressure. Then, calculate the mass of oxygen using its molar mass (approximately 32.00 g/mol) with the formula \( \text{mass} = n \times \text{molar mass} \).