Gaseous hydrocarbons, which contain only carbon and hydrogen, are good fuels because they burn in air to generate large amounts of heat. A sample of hydrocarbon with mass 1.48 g exerts a pressure of 1.32 atm in bulb with volume 967 mL at 21.5 oC. Molar Mass of this hydrocarbon is 27.9

You haven't asked a question.

To determine the number of moles of the hydrocarbon, we can use the ideal gas law, which states:

PV = nRT

Where:
P = pressure in atm
V = volume in L (1 L = 1000 mL)
n = number of moles
R = ideal gas constant (0.0821 L·atm/(mol·K))
T = temperature in Kelvin (K = °C + 273.15)

First, let's convert the given temperature to Kelvin:
T = 21.5 °C + 273.15 = 294.65 K

Next, convert the given volume to liters:
V = 967 mL ÷ 1000 mL/L = 0.967 L

Now, we can rearrange the ideal gas law to solve for the number of moles (n):
n = PV / RT

Substituting the given values into the equation:
n = (1.32 atm) * (0.967 L) / (0.0821 L·atm/(mol·K) * 294.65 K)

n = 0.0420 mol

Since the given mass of the hydrocarbon is 1.48 g, we can calculate the molar mass (M) of the hydrocarbon using the equation:

Molar mass = mass / moles

Molar mass = 1.48 g / 0.0420 mol

Molar mass = 35.2 g/mol

The calculated molar mass of the hydrocarbon is 35.2 g/mol, which does not match the given molar mass of 27.9 g/mol. Therefore, it seems that there may be an error in the given information or calculation.