Final answer:
To determine the molecular formula from an empirical formula of NO and a molar mass of 276 g/mol, the empirical formula's molar mass is calculated and compared to the given molar mass. After this comparison, one typically multiplies the subscripts of the empirical formula by the obtained ratio, but because the calculated ratio is not a whole number, our answer is questionable without additional information.
Step-by-step explanation:
To find the molecular formula of a compound that has a mass of 276 and an empirical formula of NO, you need to compare the molar mass of the empirical formula to the given molar mass of the compound. The empirical formula NO consists of Nitrogen (N) and Oxygen (O) with atomic masses of approximately 14 amu and 16 amu respectively. Therefore, the molar mass of NO is 14 + 16 = 30 g/mol.
Next, we divide the given molecular mass of the compound (276 g/mol) by the molar mass of the empirical formula (30 g/mol) to find the ratio between them:
276 ÷ 30 ≈ 9.2. Since molecular formulas are whole-number multiples of empirical formulas, we round our ratio to the nearest whole number, which is 9 in this case.
Finally, to determine the molecular formula, we multiply the subscripts in the empirical formula NO by 9, giving us a molecular formula of N9O9, which is technically incorrect because our ratio was not an exact whole number. However, compounds with NO as the empirical formula do not form such high molecular weight compounds in reality. Thus, we must infer that there may have been a mistake either in the empirical formula provided, the molar mass, or there may be a need for more information regarding the nature of the compound.