Question

The density of a test gas is to be determined experimentally at 302.2 K using an apparatus constructed of a 4.350 L glass bulb volume that is attached to a vacuum pump. The mass of the evacuated bulb is 21.072 g.?

After it is filled with the test gas to a pressure of 0.0750 atm, the mass increases to 21.310 g. Assume the gas behaves ideally. What is the density of gas? How many moles of gas are in the bulb? What is the apparent molar mass of the gas?

Answer 1

`0.0547 g/L` at those conditions.
`4.16 (g/(mol))` Probably helium.

Explanation:

Density is simply mass/volume. BUT, that is relative to the conditions for gases. Directly, the measured density at the conditions of the experiment is `(21.310 – 21.072)/4.350 = 0.0547 g/L`. Changing the pressure or temperature will change that density.

The number of moles of gas can be calculated from the Ideal Gas Laws: `PV = nRT`. With a pressure in atm and volume in L we will use an R of `0.0821 L*atm K^(-1)mol^(-1)`.

`n = (PV)/(RT)` ; `n = (0.0750 xx 4.350)/(0.0821 xx 302.2) = 0.01315` moles

The molar mass is `(mass)/("moles")`: `0.0547/0.01315 = 4.16 (g/(mol))`