Suppose that 1.00 gram of each of the three inert gases, Ne, Ar, Kr, are combined in a 1.00-L flask at 25.0°C. Determine the partial pressures of each of the three gases in the flask in torr. Ne Ar Kr Total

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To determine the partial pressures of each gas, we use the ideal gas law, $PV = nRT$, where $P$ is pressure, $V$ is volume, $n$ is the number of moles, $R$ is the ideal gas constant, and $T$ is temperature in Kelvin. The molar mass of Ne, Ar, and Kr are approximately 20.18 g/mol, 39.95 g/mol, and 83.80 g/mol, respectively. The number of moles of each gas is calculated by dividing the mass of the gas by its molar mass. The total pressure is the sum of the partial pressures of each gas. Given the conditions (1.00 L flask, 25.0°C or 298.15 K), and using $R = 62.364 \, \text{L torr K}^{-1} \text{mol}^{-1}$, the calculated partial pressures for Ne, Ar, and Kr will not be equal due to their different molar masses, leading to different amounts of moles for each gas. The correct answer must reflect the proportional differences in molar masses and the resulting partial pressures, which is accurately represented in option D, where the pressures are different for each gas and add up to a total that is plausible given the conditions.

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