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Ch.10 - Gases: Their Properties & Behavior
McMurry - Chemistry 8th Edition
McMurry8th EditionChemistryISBN: 9781292336145Not the one you use?Change textbook
All textbooksMcMurry 8th EditionCh.10 - Gases: Their Properties & BehaviorProblem 87
Chapter 10, Problem 87

A 20.0-L flask contains 0.776 g of He and 3.61 g of CO2 at300 K.(a) What is the partial pressure of He in mm Hg?

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1
Calculate the number of moles of He using the ideal gas law formula: \( n = \frac{m}{M} \), where \( m \) is the mass of the gas and \( M \) is the molar mass of He (4.00 g/mol).
Use the ideal gas law \( PV = nRT \) to find the total pressure in the flask. Here, \( P \) is the pressure, \( V \) is the volume (20.0 L), \( n \) is the total moles of gas (sum of moles of He and CO2), \( R \) is the gas constant (0.0821 L atm/mol K), and \( T \) is the temperature (300 K).
Calculate the mole fraction of He, which is the ratio of moles of He to the total moles of gases in the flask.
Determine the partial pressure of He by multiplying the total pressure by the mole fraction of He.
Convert the partial pressure of He from atm to mm Hg by using the conversion factor (1 atm = 760 mm Hg).

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Key Concepts

Here are the essential concepts you must grasp in order to answer the question correctly.

Ideal Gas Law

The Ideal Gas Law relates the pressure, volume, temperature, and number of moles of a gas through the equation PV = nRT. This law is fundamental for calculating the behavior of gases under various conditions, allowing us to determine the pressure exerted by a gas in a given volume at a specific temperature.
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Partial Pressure

Partial pressure is the pressure that a single gas in a mixture would exert if it occupied the entire volume alone at the same temperature. According to Dalton's Law of Partial Pressures, the total pressure of a gas mixture is the sum of the partial pressures of each individual gas, which is crucial for solving problems involving multiple gases.
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Molar Mass and Moles Calculation

To find the number of moles of a gas, the mass of the gas is divided by its molar mass. This calculation is essential for applying the Ideal Gas Law, as it allows us to convert the mass of the gases present (in grams) into moles, which are needed to determine the partial pressures of each gas in the mixture.
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Related Practice
Textbook Question

Natural gas is a mixture of hydrocarbons, primarily methane 1CH42 and ethane 1C2H62. A typical mixture might have Xmethane = 0.915 and Xethane = 0.085. Let's assume that we have a 15.50 g sample of natural gas in a volume of 15.00 L at a temperature of 20.00 °C. (a) How many total moles of gas are in the sample?

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Textbook Question

A mixture of 14.2 g of H2 and 36.7 g of Ar is placed in a 100.0-L container at 290 K. (b) What is the partial pressure of Ar in atmospheres?

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Textbook Question

Chlorine gas was first prepared in 1774 by the oxidation of NaCl with MnO2: 2 NaCl(s) + 2 H2SO4(l) + MnO2(s) → Na2SO4(s) + MnSO4(s) + 2 H2O(g) + Cl2(g) Assume that the gas produced is saturated with water vapor at a partial pressure of 28.7 mm Hg and that it has a volume of 0.597 L at 27 °C and 755 mm Hg pressure. (b) How many grams of NaCl were used in the experiment, assuming complete reaction?

604
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Textbook Question

A mixture of 14.2 g of H2 and 36.7 g of Ar is placed in a 100.0-L container at 290 K. (a) What is the partial pressure of H2 in atmospheres?

726
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Textbook Question

Chlorine gas was first prepared in 1774 by the oxidation of NaCl with MnO2: 2 NaCl(s) + 2 H2SO4(l) + MnO2(s) → Na2SO4(s) + MnSO4(s) + 2 H2O(g) + Cl2(g) Assume that the gas produced is saturated with water vapor at a partial pressure of 28.7 mm Hg and that it has a volume of 0.597 L at 27 °C and 755 mm Hg pressure. (a) What is the mole fraction of Cl2 in the gas?

380
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