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Chapter 3: Problem 52

What is the molecular formula of each of the following compounds? (a) empirical formula \(\mathrm{CH}_{3} \mathrm{O}\), molar mass $=62.0 \mathrm{~g} / \mathrm{mol}$ (b) empirical formula \(\mathrm{NH}_{2}\), molar mass $=32.0 \mathrm{~g} / \mathrm{mol}$

Short Answer

Expert verified
The molecular formulas of the given compounds are: (a) C2H6O2 (b) N2H4

Step by step solution

01

(a) Finding the molar mass of the empirical formula C3H3O

To find the molar mass of the empirical formula, add the atomic masses of the elements present in the formula. The atomic masses are: Carbon (C) = 12.01 g/mol Hydrogen (H) = 1.008 g/mol Oxygen (O) = 16.00 g/mol Molar mass of the empirical formula (CH3O) = 1 × 12.01 + 3 × 1.008 + 1 × 16.00 = \(12.01 + 3.024 + 16.00 = 31.034 \mathrm{~g/mol}\)
02

(a) Finding the molecular formula

Now, we will find the multiple needed to obtain the given molar mass of the compound. Divide the given molar mass by the molar mass of the empirical formula: Multiple = \(\frac{62.0\mathrm{~g/mol}}{31.034 \mathrm{~g/mol}} = 2 \) Now, multiply the empirical formula by this multiple to find the molecular formula: Molecular formula = 2 × (CH3O) = C2H6O2 So, the molecular formula for compound (a) is C2H6O2.
03

(b) Finding the molar mass of the empirical formula NH2

To find the molar mass of the empirical formula, add the atomic masses of the elements present in the formula. The atomic masses are: Nitrogen (N) = 14.01 g/mol Hydrogen (H) = 1.008 g/mol Molar mass of the empirical formula (NH2) = 1 × 14.01 + 2 × 1.008 = \(14.01 + 2.016 = 16.026 \mathrm{~g/mol}\)
04

(b) Finding the molecular formula

Now, we will find the multiple needed to obtain the given molar mass of the compound. Divide the given molar mass by the molar mass of the empirical formula: Multiple = \(\frac{32.0\mathrm{~g/mol}}{16.026\mathrm{~g/mol}} = 2 \) Now, multiply the empirical formula by this multiple to find the molecular formula: Molecular formula = 2 × (NH2) = N2H4 So, the molecular formula for compound (b) is N2H4.

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Most popular questions from this chapter

When ethane \(\left(\mathrm{C}_{2} \mathrm{H}_{6}\right)\) reacts with chlorine \(\left(\mathrm{Cl}_{2}\right)\), the main product is $\mathrm{C}_{2} \mathrm{H}_{5} \mathrm{Cl}\(, but other products containing \)\mathrm{Cl}$, such as \(\mathrm{C}_{2} \mathrm{H}_{4} \mathrm{Cl}_{2}\), are also obtained in small quantities. The formation of these other products reduces the yield of \(\mathrm{C}_{2} \mathrm{H}_{5} \mathrm{Cl}\). (a) Calculate the theoretical yield of $\mathrm{C}_{2} \mathrm{H}_{5} \mathrm{Cl}\( when \)125 \mathrm{~g}\( of \)\mathrm{C}_{2} \mathrm{H}_{6}$ reacts with \(255 \mathrm{~g}\) of \(\mathrm{Cl}_{2}\), assuming that $\mathrm{C}_{2} \mathrm{H}_{6}\( and \)\mathrm{Cl}_{2}\( react only to form \)\mathrm{C}_{2} \mathrm{H}_{2} \mathrm{Cl}$ and HCl. (b) Calculate the percent yield of \(\mathrm{C}_{2} \mathrm{H}_{5} \mathrm{Cl}\) if the reaction produces $206 \mathrm{~g}\( of \)\mathrm{C}_{2} \mathrm{H}_{5} \mathrm{Cl}$.

A compound whose empirical formula is \(\mathrm{XF}_{3}\) consists of \(65 \%\) \(\mathrm{F}\) by mass. What is the atomic mass of \(\mathrm{X} ?\)

Balance the following equations: (a) $\mathrm{CaS}(s)+\mathrm{H}_{2} \mathrm{O}(l) \longrightarrow \mathrm{Ca}(\mathrm{HS})_{2}(a q)+\mathrm{Ca}(\mathrm{OH})_{2}(a q)$ (b) $\mathrm{NH}_{3}(g)+\mathrm{O}_{2}(g) \longrightarrow \mathrm{NO}(g)+\mathrm{H}_{2} \mathrm{O}(g)$ (c) $\mathrm{FeCl}_{3}(s)+\mathrm{Na}_{2} \mathrm{CO}_{3}(a q) \longrightarrow \mathrm{Fe}_{2}\left(\mathrm{CO}_{3}\right)_{3}(s)+\mathrm{NaCl}(a q)$ (d) $\mathrm{FeS}_{2}(s)+\mathrm{O}_{2}(g) \longrightarrow \mathrm{Fe}_{2} \mathrm{O}_{3}(s)+\mathrm{SO}_{2}(g)$

A piece of aluminum foil \(1.00 \mathrm{~cm}^{2}\) and \(0.550-\mathrm{mm}\) thick is allowed to react with bromine to form aluminum bromide. (a) How many moles of aluminum were used? (The density of aluminum is $2.699 \mathrm{~g} / \mathrm{cm}^{3} .$ ) (b) How many grams of aluminum bromide form, assuming the aluminum reacts completely?

Write balanced chemical equations for \((\mathbf{a})\) the complete combustion of acetone \(\left(\mathrm{CH}_{3} \mathrm{COCH}_{3}\right),\) a common organic solvent; (b) the decomposition of solid mercury (I) carbonate into carbon dioxide gas, mercury, and solid mercury oxide; (c) the combination reaction between sulphur dioxide gas and liquid water to produce sulfurous acid.
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