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To identify a diatomic gas (X2), a researcher carried out the following experiment: She weighed an empty 2.2-L bulb, then filled it with the gas at 1.70 atm and 26.0 āˆ˜C and weighed it again. The difference in mass was 4.3 g . Identify the gas.

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Answer:

The gas is Nā‚‚

Explanation:

To solve this problem we need to use PV=nRT. The problem gives us V, P and T, so first let's convert 26.0 Ā°C into K:

26.0 + 273.16 = 299.16 K

There's no need to convert V and P as they already are in proper units.

  • Calculating n:

PV = nRT

1.70 atm * 2.2 L = n * 0.082 atmĀ·LĀ·molā»Ā¹Ā·Kā»Ā¹ * 299.16 K

n = 0.1524 mol

From the difference in mass of the experiment, we know that the mass of the gass is 4.3 g. We also now know that 0.1524 moles of Xā‚‚ weigh 4.3 g. Now we calculate the molecular weight:

4.3 g / 0.1524 mol = 28.22 g/mol

  • Two atoms of X weigh 28.22, so one atom weighs 14.11

Looking at the periodic table, the element with a similar atomic weight is N. So the gas is Nā‚‚.