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In this lesson, we will learn how to use the ideal gas law to calculate the molar mass of an ideal gas from its density, temperature, and pressure.

Q1:

The density of a certain gaseous fluoride of phosphorus is 3.88 g/L at standard temperature and pressure ( 0 . 0 ∘ C , 1.00 bar).

What is the molar mass of this fluoride?

What is the molecular formula of this fluoride?

Q2:

What is the molar mass of an ideal gas if 0.281 g of the gas occupies a volume of 125 mL at a temperature 1 2 6 ∘ C and a pressure of 777 torr?

Q3:

The approximate molar mass of a volatile liquid can be determined by:

What of the following assumptions is not necessary in order to calculate a good estimate of the molar mass from the mass of the vapor?

Using this procedure a sample of chloroform gas weighing 0.494 g is collected in a flask with a volume of 129 cm^{3} at 9 9 . 6 ∘ C when the atmospheric pressure is 742.1 mmHg. What is the approximate molar mass of chloroform?

Q4:

At a temperature of 5 0 0 ∘ C and a pressure of 0.932 bar, sulfur vapor has a density of 3.71 g/dm^{3}. What molecular formula for sulfur is compatible with this set of conditions?

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