What is the mass of 1.23 L of methane gas at STP?

What is the molar volume for a gas at -70.2^{o}C and 4183.8 mmHg?

An imaginary ideal gas has a density of 3.301 g/L at STP. What is the molar mas of this gas

1st Question)

Use the ideal gas equation: PV = nRT

Isolate Moles (n) in the equation: n = PV/RT

At STP, Pressure is 1 atm Temperature is 273 K, and R = 0.08206

Plug in everything to n = PV/RT

n = (1 atm * 1.23 L) / (0.08206 * 273 k)

n =0.0549 moles of CH4 (g)

Multiply 0.0549 moles by 16.04 gram/mole to get 0.8807 grams of CH4 (g) at STP

2nd Question)

Isolate Volume in the equation: V = nRT / P

Convert Celsius to Kelvin: -70.2 C = 202.8 K

Convert mmHg to Atm: 4183.8 mmHg = 5.505 atm

Plug in everything to V = nRT / P

V = (1 mole * 0.08206 * 202.8 K) / (5.505 atm)

Molar Volume at those conditions = 3.023 L

3rd Question)

Use the modified ideal gas equation: Density = (Pressure * Molar mass) / (R * Temperature)

Isolate the Molar mass = (Density * Rate * Temperature) / Pressure

Plug in everything to Molar mass = (Density * Rate * Temperature) / Pressure

Molar mass = (3.301*0.08206*273) / (1 atm) = 73.95 grams / mole

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