91Ó°ÊÓ

Consider the following balanced chemical equation: \(2 \mathrm{H}_{2}+\mathrm{O}_{2} \rightarrow 2 \mathrm{H}_{2} \mathrm{O}\) (a) How many grams of water are formed from \(5.00 \mathrm{~g}\) of \(\mathrm{H}_{2}\) and an excess amount of \(\mathrm{O}_{2}\) ? (b) How many grams of \(\mathrm{O}_{2}\) do you need to produce \(5.00 \mathrm{~g}\) of \(\mathrm{H}_{2} \mathrm{O}\) ? (c) Given \(100.0 \mathrm{~g}\) of \(\mathrm{H}_{2}\), how many grams of \(\mathrm{O}_{2}\) are required to run the reaction in a stoichiometric fashion? (d) What is the theoretical yield in grams of water upon combining \(50.0 \mathrm{~g}\) of \(\mathrm{O}_{2}\) with an excess amount of \(\mathrm{H}_{2}\) ? (e) Express the answer to part (d) in terms of the number of water molecules.

Chlorine \(\left(\mathrm{Cl}_{2}\right)\) and fluorine \(\left(\mathrm{F}_{2}\right)\) react to form \(\mathrm{ClF}_{3}\). A reaction vessel contains \(2.50\) moles of \(\mathrm{Cl}_{2}\) and \(6.15\) moles of \(\mathrm{F}_{2}\). (a) Write a balanced chemical equation for this reaction. (b) Which reactant is limiting?

Chlorine \(\left(\mathrm{Cl}_{2}\right)\) and fluorine \(\left(\mathrm{F}_{2}\right)\) react to form \(\mathrm{ClF}_{3}\). A reaction vessel contains \(10.00 \mathrm{~g} \mathrm{Cl}_{2}\) and \(10.00 \mathrm{~g} \mathrm{~F}_{2}\). (Hint: Refer to Problem 9.82.) (a) Write a balanced chemical equation for this reaction. (b) Which reactant is limiting? (c) What is the theoretical yield for this reaction in grams? (d) How many grams of excess reactant are left over at the end of the reaction? (e) When this reaction is actually performed, \(12.50 \mathrm{~g}\) of \(\mathrm{ClF}_{3}\) is recovered. What is the percent yield of the reaction?

Sodium (Na) reacts with hydrogen \(\left(\mathrm{H}_{2}\right)\) to form sodium hydride (NaH). A reaction mixture contains \(10.00 \mathrm{~g} \mathrm{Na}\) and \(0.0235 \mathrm{~g} \mathrm{H}_{2}\). (a) Write a balanced chemical equation for this reaction. (b) Which reactant is limiting? (c) What is the theoretical yield for this reaction in grams? (d) How many grams of excess reactant are left over at the end of the reaction? (e) When this reaction is actually performed, \(0.428 \mathrm{~g}\) of \(\mathrm{NaH}\) is recovered. What is the percent yield of the reaction?

An organic compound of carbon and hydrogen has the empirical formula CH. What is its molecular formula if its molar mass is: (a) \(26 \mathrm{~g} / \mathrm{mol}\) (b) \(52 \mathrm{~g} / \mathrm{mol}\) (c) \(78 \mathrm{~g} / \mathrm{mol}\)

A \(1.540-g\) sample of a liquid is subjected to combustion analysis, yielding \(40.00 \% \mathrm{C}\) and \(6.71 \% \mathrm{H}\). It may also contain oxygen. (a) What is the empirical formula for this compound? (b) The molar mass of this compound is determined to be about \(30 \mathrm{~g} / \mathrm{mol}\). What is the molecular formula for this compound?

Ethanol, the alcohol in beer and wine, has the molecular formula \(\mathrm{C}_{2} \mathrm{H}_{6} \mathrm{O} .\) Calculate the mass percent of each element in ethanol.

Ethylene glycol, used for antifreeze, has the molecular formula \(\mathrm{C}_{2} \mathrm{H}_{6} \mathrm{O}_{2}\). Calculate the mass percent of each element in ethylene glycol.

The compound \(\mathrm{P}_{4} \mathrm{O}_{10}\) has an empirical formula of \(\mathrm{P}_{2} \mathrm{O}_{5} .\) By what factor will the percent by mass composition differ for each element between these two formulas? Explain your answer.

The thyroid hormone thyroxine has the molecular formula \(\mathrm{C}_{15} \mathrm{H}_{11} \mathrm{NO}_{4} \mathrm{I}_{4} .\) Calculate the mass percent of each element in thyroxine.