To identify a diatomic gas (), a researcher carried out the following experiment: She weighed an empty 1.00- bulb, then filled it with the gas at 1.10 and 23.0 and weighed it again. The difference in mass was 1.27 . Identify the gas.
Express your answer as a chemical formula.
Refer to my previous post. You need units.
To identify the gas, we need to use the information provided - the mass of the gas and the conditions under which it was weighed.
First, let's calculate the mass of the gas. We are given that the difference in mass of the bulb before and after filling it with gas is 1.27 g.
Now, we need to convert this mass difference into moles. To do this, we need to know the molar mass of the unknown gas. However, we are not given this information.
To proceed, we can make use of the ideal gas law, which states that PV = nRT, where P is the pressure, V is the volume, n is the number of moles, R is the ideal gas constant, and T is the temperature in Kelvin.
In this case, we are given the pressure (1.10 atm) and temperature (23.0 °C), which we need to convert to Kelvin by adding 273.15.
Now, we can rearrange the ideal gas law to solve for n:
n = PV / RT
Using the given values:
n = (1.10 atm) * (1.00 L) / (0.0821 L·atm/mol·K * (23.0 + 273.15) K)
Simplifying the equation gives:
n = 1.10 / (0.0821 * 296.15)
n ≈ 0.047 mol
Since the gas is diatomic, each molecule contains 2 moles of atoms. Therefore, the number of molecules is twice the number of moles.
Number of molecules ≈ 2 * 0.047 mol
Now, comparing the number of molecules to the mass difference:
2 * 0.047 mol ≈ 1.27 g
Dividing the mass difference by the number of molecules to find the molar mass:
(1.27 g) / (2 * 0.047 mol) ≈ 13.5 g/mol
At this point, we need to consult the periodic table to determine the most likely diatomic gas with a molar mass of approximately 13.5 g/mol. Oxygen (O2) is the most likely candidate, as its molar mass is close to the calculated value.
Therefore, the gas is most likely diatomic oxygen, O2.