Final answer:
By using the Ideal Gas Law and the information provided, the calculated molar mass of the gas is approximately 44.3 g/mol, which does not correspond to any of the options given. This suggests a possible discrepancy in the options or the experimental data.
Step-by-step explanation:
To find the molar mass of the gas collected in the general chemistry laboratory experiment, we'll use the Ideal Gas Law, which can be expressed as 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.
First, convert the given temperature to Kelvin: T = 27 °C + 273.15 = 300.15 K.
Then, we have to adjust the pressure to account for the vapor pressure of the water. However, since the vapor pressure of water at 27 °C is not provided, we'll assume it's negligible for the purpose of this calculation and use the pressure as given (this is a simplification and in practice the vapor pressure should be subtracted from the total pressure).
The volume needs to be converted to liters: V = 265 mL = 0.265 L.
The pressure is given in torr and should be converted to atm for the Ideal Gas Constant (R) in L atm/(mol K): P = 753 torr * (1 atm / 760 torr) ≈ 0.991 atm.
We will use the Ideal Gas Constant R = 0.0821 L atm/(mol K).
Now we can solve for the number of moles (n):
n = PV / RT = (0.991 atm * 0.265 L) / (0.0821 L atm/(mol K) * 300.15 K) ≈ 0.01065 mol.
Finally, calculate the molar mass (M) using the mass (m) of the gas and the number of moles (n): M = m / n = 0.472 g / 0.01065 mol ≈ 44.3 g/mol.
As the calculated molar mass does not match any of the provided options (a) 17.2 g/mol, (b) 26.8 g/mol, (c) 31.5 g/mol, (d) 40.6 g/mol, it is possible there is an error in either the provided options or the experiment itself.