Question

To identify a diatomic gas (X2), a researcher carried out the following experiment: She weighed an empty 4.6-L bulb, then filled it with the gas at 1.80 atm and 27.0 ∘C and weighed it again. The difference in mass was 9.5 g . Identify the gas. Express your answer as a chemical formula.

Answer 1

Answer:

The diatomic gas is probably N2

Explanation:

Step 1: Data given

Volume of the bulb = 4.6 L

Pressure = 1.80 atm

Temperature = 27.0 °C = 300K

Mass of the gas = 9.5 grams

Step 2: Calculate the number of moles

PV=nRT

⇒with P = the pressure of the gas= 1.80 atm

⇒with V = the volume of the gas = 4.6 L

⇒with n = the number of moles gas = TO BE DETERMINED

⇒with R = the gas constant = 0.08206 L*atm/mol*K

⇒with T = the temperature = 300 K

n = (p*V)/(R*T)

n = (1.80 * 4.6) / (0.08206*300)

n = 0.336 moles

Step 3: Calculate molar mass

Molar mass = mass / moles

Molar mass = 9.5 grams / 0.336 moles

Molar mass = 28.3 g/mol

The diatomic gas is probably N2