To identify a diatomic gas (), a researcher carried out the following experiment: She weighed an empty 1.00- bulb, then filled it with the gas at 1.10 and 23.0 and weighed it again. The difference in mass was 1.27 . Identify the gas.
The problem is not solvable without units. Please recopy the problem and put in the units of the problem.
Express your answer as a chemical formula.
not solvable
To identify the gas, we need to determine the molar mass of the gas based on the given experimental data. We can then compare this molar mass to the molar masses of various diatomic gases to identify the gas in question.
1. Start by calculating the number of moles of the gas used in the experiment using the ideal gas law equation:
PV = nRT
Where:
P = pressure (1.10 atm)
V = volume of the bulb (1.00 L)
n = number of moles (to be determined)
R = ideal gas constant (0.0821 L·atm/(mol·K))
T = temperature in Kelvin (23.0 + 273.15 K)
Rearranging the equation to solve for n:
n = PV / (RT)
Substituting the given values:
n = (1.10 atm * 1.00 L) / (0.0821 L·atm/(mol·K) * (23.0 + 273.15 K))
Calculate the value of 'n'.
2. Next, determine the molar mass of the gas by dividing the change in mass of the bulb (1.27 g) by the moles of gas calculated in step 1.
Molar mass = change in mass / number of moles
Substitute the given values:
Molar mass = 1.27 g / calculated value of 'n'
Calculate the molar mass.
3. Compare the calculated molar mass to the molar masses of common diatomic gases to identify the gas. Some common diatomic gases and their molar masses are:
- Molecular nitrogen (N2): approximately 28.01 g/mol
- Molecular oxygen (O2): approximately 32.00 g/mol
- Hydrogen gas (H2): approximately 2.02 g/mol
Compare the calculated molar mass to these possible options to determine the gas in question.
By following these steps, you should be able to identify the diatomic gas based on the given experimental data.