Question

Nitrogen forms several gaseous oxides. One of them has a density of 1.33 g/l . measured at 764 mm Hg and ${150}^{\circ }C.$ Write the formula of the compound.

• A. $NO$
• B. $N_{2}O$
• C. $N{O}_2$
• D. $N_2{O}_s$

Answer 1

The correct option is C $N{O}_2$

1. Use the Ideal Gas Equation: PV = nRT and derive the equation for molar mass
PV = (mass/molar mass) RT
P * molar mass = mass/V * RT
but mass/V is equal to density
So,
Molar Mass = (D * R * T)/ P
2. Convert the Temperature to Kelvin
T = 150 + 273 = 423 K
3. Convert 764 mm Hg to atmosphere
P = 764 mm * (1 atm/760 mm) = 1.01 atm
4. Solve for the molar mass using the information from steps 2 and 3 and the derived equation in step 1.
Molar Mass =[(1.33 g/L)*(0.08206 L.atm/mol.K)* 423K] / 1.01 atm
Molar Mass = 45.7 = 46 g/mol
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The possible compounds with low molar masses are NO, N2O and NO2.
Calculate the molar masses and compare with the molar mass from the given info.
NO = 14.01 + 16.00 = 30.01 g/mol --> NOT the compound
N2O = 2*14.01 + 16.00 = 44.01 g/mol --> NOT the compound
NO2 = 14.01 + 2*16.00 = 46.01 g/mol --> This is the MATCH!