To find the molar mass ($M$) of the compound, we can use the Ideal Gas Law:

$PV = nRT$

where:

• $P$ = pressure of the gas (1.00 atm),
• $V$ = volume of the gas (2,100 mL = 2.1 L),
• $n$ = number of moles of the gas,
• $R$ = ideal gas constant (0.0821 L·atm / K·mol),
• $T$ = temperature in Kelvin (140°C + 273.15 = 413.15 K).

1. Calculate the moles of gas ($n$):

   Rearrange the Ideal Gas Law to solve for $n$: $n = \frac{PV}{RT}$

   Substitute the values: $n = \frac{(1.00 \, \text{atm})(2.1 \, \text{L})}{(0.0821 \, \text{L·atm / K·mol})(413.15 \, \text{K})}$

   Calculate $n$: $n \approx \frac{2.1}{33.916215} \approx 0.0619 \, \text{moles}$

2. Calculate the molar mass ($M$):

   Molar mass ($M$) is defined as the mass of the compound divided by the number of moles: $M = \frac{\text{mass}}{n}$

   Substitute the values: $M = \frac{1.99 \, \text{g}}{0.0619 \, \text{moles}}$

   Calculate $M$: $M \approx 32.1 \, \text{g/mol}$

Answer: The molar mass of the compound is approximately $32.1 \, \text{g/mol}$, rounded to three significant figures.

Would you like more details on any part of this calculation?

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Here are some related questions to explore:

1. How would changing the temperature affect the calculated molar mass?
2. What assumptions do we make when using the Ideal Gas Law for real gases?
3. How do significant figures impact the accuracy of scientific calculations?
4. What are the consequences if the gas does not behave ideally under these conditions?
5. How does the choice of units for pressure and volume affect the Ideal Gas Law calculation?

Tip: Always convert temperatures to Kelvin when using the Ideal Gas Law; the formula requires absolute temperature to ensure accuracy.