Chapter 4

Write a balanced chemical equation for the combustion of liquid pentane.

Write a balanced chemical equation for the production of ammonia, \(\mathrm{NH}_{3}(\mathrm{g}),\) from \(\mathrm{N}_{2}(\mathrm{g})\) and \(\mathrm{H}_{2}(\mathrm{g})\)

Balance the following equations: (a) \(\operatorname{Cr}(\mathrm{s})+\mathrm{O}_{2}(\mathrm{g}) \longrightarrow \mathrm{Cr}_{2} \mathrm{O}_{3}(\mathrm{s})\) (b) \(\mathrm{Cu}_{2} \mathrm{S}(\mathrm{s})+\mathrm{O}_{2}(\mathrm{g}) \longrightarrow \mathrm{Cu}(\mathrm{s})+\mathrm{SO}_{2}(\mathrm{g})\) (c) \(\mathrm{C}_{6} \mathrm{H}_{5} \mathrm{CH}_{3}(\ell)+\mathrm{O}_{2}(\mathrm{g}) \longrightarrow \mathrm{H}_{2} \mathrm{O}(\ell)+\mathrm{CO}_{2}(\mathrm{g})\)

Balance the following equations: (a) \(\mathrm{Cr}(\mathrm{s})+\mathrm{Cl}_{2}(\mathrm{g}) \longrightarrow \mathrm{Cr} \mathrm{Cl}_{3}(\mathrm{s})\) (b) \(\mathrm{SiO}_{2}(\mathrm{s})+\mathrm{C}(\mathrm{s}) \longrightarrow \mathrm{Si}(\mathrm{s})+\mathrm{CO}(\mathrm{g})\) (c) \(\mathrm{Fe}(\mathrm{s})+\mathrm{H}_{2} \mathrm{O}(\mathrm{g}) \longrightarrow \mathrm{Fe}_{3} \mathrm{O}_{4}(\mathrm{s})+\mathrm{H}_{2}(\mathrm{g})\)

Balance the following equations and name each reactant and product: (a) \(\mathrm{Fe}_{2} \mathrm{O}_{3}(\mathrm{s})+\mathrm{Mg}(\mathrm{s}) \longrightarrow \mathrm{MgO}(\mathrm{s})+\mathrm{Fe}(\mathrm{s})\) (b) \(\mathrm{AlCl}_{3}(\mathrm{s})+\mathrm{NaOH}(\mathrm{aq}) \longrightarrow \mathrm{Al}(\mathrm{OH})_{3}(\mathrm{s})+\mathrm{NaCl}(\mathrm{aq})\) (c) \(\mathrm{NaNO}_{3}(\mathrm{s})+\mathrm{H}_{2} \mathrm{SO}_{4}(\ell) \longrightarrow \mathrm{Na}_{2} \mathrm{SO}_{4}(\mathrm{s})+\mathrm{HNO}_{3}(\ell)\) (d) \(\mathrm{NiCO}_{3}(\mathrm{s})+\mathrm{HNO}_{3}(\mathrm{aq}) \longrightarrow\) \(\mathrm{Ni}\left(\mathrm{NO}_{3}\right)_{2}(\mathrm{aq})+\mathrm{CO}_{2}(\mathrm{g})+\mathrm{H}_{2} \mathrm{O}(\ell)\)

Balance the following equations and name each reactant and product: (a) \(\mathrm{SF}_{4}(\mathrm{g})+\mathrm{H}_{2} \mathrm{O}(\ell) \longrightarrow \mathrm{SO}_{2}(\mathrm{g})+\mathrm{HF}(\ell)\) (b) \(\mathrm{NH}_{3}(\mathrm{aq})+\mathrm{O}_{2}(\mathrm{aq}) \longrightarrow \mathrm{NO}(\mathrm{g})+\mathrm{H}_{2} \mathrm{O}(\ell)\) (c) \(\mathrm{BF}_{3}(\mathrm{g})+\mathrm{H}_{2} \mathrm{O}(\ell) \longrightarrow \mathrm{HF}(\mathrm{aq})+\mathrm{H}_{3} \mathrm{BO}_{3}(\mathrm{aq})\)

Aluminum reacts with oxygen to give aluminum oxide. $$ 4 \mathrm{Al}(\mathrm{s})+3 \mathrm{O}_{2}(\mathrm{g}) \longrightarrow 2 \mathrm{Al}_{2} \mathrm{O}_{3}(\mathrm{s}) $$ What amount of \(\mathrm{O}_{2},\) in moles, is needed for complete reaction with 6.0 mol of Al? What mass of \(\mathrm{Al}_{2} \mathrm{O}_{3},\) in grams, can be produced?

What mass of HCl, in grams, is required to react with \(0.750 \mathrm{g}\) of \(\mathrm{Al}(\mathrm{OH})_{3} ?\) What mass of water, in grams, is produced? \(\mathrm{Al}(\mathrm{OH})_{3}(\mathrm{s})+3 \mathrm{HCl}(\mathrm{aq}) \longrightarrow \mathrm{AlCl}_{3}(\mathrm{aq})+3 \mathrm{H}_{2} \mathrm{O}(\ell)\)

Like many metals, aluminum reacts with a halogen to give a metal halide (see Figure 3.1 ). $$ 2 \mathrm{Al}(\mathrm{s})+3 \mathrm{Br}_{2}(\ell) \longrightarrow \mathrm{Al}_{2} \mathrm{Br}_{6}(\mathrm{s}) $$ What mass of \(\mathrm{Br}_{2}\), in grams, is required for complete reaction with \(2.56 \mathrm{g}\) of Al? What mass of white, solid \(\mathrm{Al}_{2} \mathrm{Br}_{6}\) is expected?

The balanced equation for a reaction in the process of reducing iron ore to the metal is $$ \mathrm{Fe}_{2} \mathrm{O}_{3}(\mathrm{s})+3 \mathrm{CO}(\mathrm{g}) \longrightarrow 2 \mathrm{Fe}(\mathrm{s})+3 \mathrm{CO}_{2}(\mathrm{g}) $$ (a) What is the maximum mass of iron, in grams, that can be obtained from \(454 \mathrm{g}(1.00 \mathrm{lb})\) of iron(III) oxide? (b) What mass of \(\mathrm{CO}\) is required to react with \(454 \mathrm{g}\) of \(\mathrm{Fe}_{2} \mathrm{O}_{3} ?\)

Iron metal reacts with oxygen to give iron(III) oxide, \(\mathrm{Fe}_{2} \mathrm{O}_{3}\) (a) Write a balanced equation for the reaction. (b) If an ordinary iron nail (assumed to be pure iron) has a mass of \(2.68 \mathrm{g},\) what mass of \(\mathrm{Fe}_{2} \mathrm{O}_{3},\) in grams, is produced if the nail is converted completely to the oxide? (c) What mass of \(\mathrm{O}_{2}\), in grams, is required for the reaction?

Methane, \(\mathrm{CH}_{4},\) burns in oxygen. (a) What are the products of the reaction? (b) In Write the balanced equation for the reaction. (c) What mass of \(\mathrm{O}_{2}\), in grams, is required for complete combustion of \(25.5 \mathrm{g}\) of methane? (d) What is the total mass of products expected from the combustion of \(25.5 \mathrm{g}\) of methane?

Sulfur dioxide, a pollutant produced by burning coal and oil in power plants, can be removed by reaction with calcium carbonate. \(2 \mathrm{SO}_{2}(\mathrm{g})+2 \mathrm{CaCO}_{3}(\mathrm{s})+\mathrm{O}_{2}(\mathrm{g}) \longrightarrow 2 \mathrm{CaSO}_{4}(\mathrm{s})+2 \mathrm{CO}_{2}(\mathrm{g})\) (a) What mass of \(\mathrm{CaCO}_{3}\) is required to remove \(155 \mathrm{g}\) of \(\mathrm{SO}_{2} ?\) (b) What mass of \(\mathrm{CaSO}_{4}\) is formed when \(155 \mathrm{g}\) of \(\mathrm{SO}_{2}\) is consumed completely?

The formation of water-insoluble silver chloride is useful in the analysis of chloride-containing substances. Consider the following unbalanced equation: \(\mathrm{BaCl}_{2}(\mathrm{aq})+\mathrm{AgNO}_{3}(\mathrm{aq}) \longrightarrow \mathrm{AgCl}(\mathrm{s})+\mathrm{Ba}\left(\mathrm{NO}_{3}\right)_{2}(\mathrm{aq})\) (a) Write the balanced equation. (b) What mass AgNO \(_{3},\) in grams, is required for complete reaction with \(0.156 \mathrm{g}\) of \(\mathrm{BaCl}_{2} ?\) What mass of \(\mathrm{AgCl}\) is produced?

A major source of air pollution years ago was the metals industry. One common process involved "roasting" metal sulfides in the air: $$ 2 \mathrm{PbS}(\mathrm{s})+3 \mathrm{O}_{2}(\mathrm{g}) \longrightarrow 2 \mathrm{PbO}(\mathrm{s})+2 \mathrm{SO}_{2}(\mathrm{g}) $$ If you heat 2.5 mol of \(\mathrm{PbS}\) in the air, what amount of \(\mathrm{O}_{2}\) is required for complete reaction? What amounts of \(\mathrm{PbO}\) and \(\mathrm{SO}_{2}\) are expected?

Iron ore is converted to iron metal in a reaction with carbon. $$ 2 \mathrm{Fe}_{2} \mathrm{O}_{3}(\mathrm{s})+3 \mathrm{C}(\mathrm{s}) \longrightarrow 4 \mathrm{Fe}(\mathrm{s})+3 \mathrm{CO}_{2}(\mathrm{g}) $$ If 6.2 mol of \(\mathrm{Fe}_{2} \mathrm{O}_{3}(\mathrm{s})\) is used, what amount of \(\mathrm{C}(\mathrm{s})\) is needed and what amounts of Fe and \(\mathrm{CO}_{2}\) are produced?

Chromium metal reacts with oxygen to give chromium(III) oxide, \(\mathrm{Cr}_{2} \mathrm{O}_{3}\) (a) Write a balanced equation for the reaction. (b) If a piece of chromium has a mass of \(0.175 \mathrm{g},\) what mass (in grams) of \(\mathrm{Cr}_{2} \mathrm{O}_{3}\) is produced if the metal is converted completely to the oxide? (c) What mass of \(\mathrm{O}_{2}\) (in grams) is required for the reaction?

Ethane, \(\mathrm{C}_{2} \mathrm{H}_{6},\) burns in oxygen. (a) What are the products of the reaction? (b) Write the balanced equation for the reaction. (c) What mass of \(\mathrm{O}_{2}\), in grams, is required for complete combustion of 13.6 of ethane? (d) What is the total mass of products expected from the combustion of \(13.6 \mathrm{g}\) of ethane?

Sodium sulfide, \(\mathrm{Na}_{2} \mathrm{S}\), is used in the leather industry to remove hair from hides. (This is the reason these kinds of plants stink!) The \(\mathrm{Na}_{2} \mathrm{S}\) is made by the reaction \(\mathrm{Na}_{2} \mathrm{SO}_{4}(\mathrm{s})+4 \mathrm{C}(\mathrm{s}) \longrightarrow \mathrm{Na}_{2} \mathrm{S}(\mathrm{s})+4 \mathrm{CO (\mathrm{g})\) Suppose you mix \(15 \mathrm{g}\) of \(\mathrm{Na}_{2} \mathrm{SO}_{4}\) and \(7.5 \mathrm{g}\) of \(\mathrm{C} .\) Which is the limiting reactant? What mass of \(\mathrm{Na}_{2} \mathrm{S}\) is produced?

The compound \(\mathrm{SF}_{6}\) is made by burning sulfur in an atmosphere of fluorine. The balanced equation is $$ \mathrm{S}_{8}(\mathrm{~s})+24 \mathrm{~F}_{2}(\mathrm{~g}) \longrightarrow 8 \mathrm{SF}_{6}(\mathrm{~g}) $$If you begin with 1.6 moles of sulfur, \(\mathrm{S}_{8}\), and 35 moles of \(\mathrm{F}_{2}\), which is the limiting reagent?

The reaction of methane and water is one way to prepare hydrogen for use as a fuel: $$ \mathrm{CH}_{4}(\mathrm{g})+\mathrm{H}_{2} \mathrm{O}(\mathrm{g}) \longrightarrow \mathrm{CO}(\mathrm{g})+3 \mathrm{H}_{2}(\mathrm{g}) $$ If you begin with \(995 \mathrm{g}\) of \(\mathrm{CH}_{4}\) and \(2510 \mathrm{g}\) of water, (a) Which reactant is the limiting reactant? (b) What is the maximum mass of \(\mathrm{H}_{2}\) that can be prepared? (c) What mass of the excess reactant remains when the reaction is completed?

Aluminum chloride, AlCl_, is made by treating scrap eluminum with chlorine. $$ 2 \mathrm{Al}(\mathrm{s})+3 \mathrm{Cl}_{2}(\mathrm{g}) \longrightarrow 2 \mathrm{AlCl}_{3}(\mathrm{s}) $$ If you begin with \(2.70 \mathrm{g}\) of \(\mathrm{Al}\) and \(4.05 \mathrm{g}\) of \(\mathrm{Cl}_{2}\) (a) Which reactant is limiting? (b) What mass of AlCl \(_{3}\) can be produced? (c) What mass of the excess reactant remains when the reaction is completed?

Hexane \(\left(\mathrm{C}_{6} \mathrm{H}_{14}\right)\) burns in air \(\left(\mathrm{O}_{2}\right)\) to give \(\mathrm{CO}_{2}\) and \(\mathrm{H}_{2} \mathrm{O}\) (a) Write a balanced equation for the reaction. (b) If \(215 \mathrm{g}\) of \(\mathrm{C}_{6} \mathrm{H}_{14}\) is mixed with \(215 \mathrm{g}\) of \(\mathrm{O}_{2},\) what masses of \(\mathrm{CO}_{2}\) and \(\mathrm{H}_{2} \mathrm{O}\) are produced in the reaction? (c) What mass of the excess reactant remains after the hexane has been burned?

In Example 4.3 you found that a mixture of \(\mathrm{CO}\) and \(\mathrm{H}_{2}\) produced \(407 \mathrm{g} \mathrm{CH}_{3} \mathrm{OH}\) $$ \mathrm{CO}(\mathrm{g})+2 \mathrm{H}_{2}(\mathrm{g}) \longrightarrow \mathrm{CH}_{3} \mathrm{OH}(\ell) $$ If only \(332 \mathrm{g}\) of \(\mathrm{CH}_{3} \mathrm{OH}\) is actually produced, what is the percent yield of the compound?

Ammonia gas can be prepared by the following reaction: \(\mathrm{CaO}(\mathrm{s})+2 \mathrm{NH}_{4} \mathrm{Cl}(\mathrm{s}) \longrightarrow 2 \mathrm{NH}_{3}(\mathrm{g})+\mathrm{H}_{2} \mathrm{O}(\mathrm{g})+\mathrm{CaCl}_{2}(\mathrm{s})\) If \(112 \mathrm{g}\) of \(\mathrm{CaO}\) and \(224 \mathrm{g}\) of \(\mathrm{NH}_{4} \mathrm{Cl}\) are mixed, the theoretical yield of \(\mathrm{NH}_{3}\) is \(68.0 \mathrm{g}\) (Study Question 20 ). If only \(16.3 \mathrm{g}\) of \(\mathrm{NH}_{3}\) is actually obtained, what is its percent yield?

The deep blue compound \(\mathrm{Cu}\left(\mathrm{NH}_{3}\right)_{4} \mathrm{SO}_{4}\) is made by the reaction of copper(II) sulfate and ammonia. \(\mathrm{CuSO}_{4}(\mathrm{aq})+4 \mathrm{NH}_{3}(\mathrm{aq}) \longrightarrow \mathrm{Cu}\left(\mathrm{NH}_{3}\right)_{4} \mathrm{SO}_{4}(\mathrm{aq})\) (a) If you use \(10.0 \mathrm{g}\) of \(\mathrm{CuSO}_{4}\) and excess \(\mathrm{NH}_{3}\), what is the theoretical yield of \(\mathrm{Cu}\left(\mathrm{NH}_{3}\right)_{4} \mathrm{SO}_{4} ?\) (b) If you isolate \(12.6 \mathrm{g}\) of \(\mathrm{Cu}\left(\mathrm{NH}_{3}\right)_{4} \mathrm{SO}_{4},\) what is the percent yield of \(\mathrm{Cu}\left(\mathrm{NH}_{3}\right)_{4} \mathrm{SO}_{4} ?\)

A reaction studied by Wächtershäuser and Huber (see "Black Smokers and the Origins of Life") is $$ 2 \mathrm{CH}_{3} \mathrm{SH}+\mathrm{CO} \longrightarrow \mathrm{CH}_{3} \mathrm{COSCH}_{3}+\mathrm{H}_{2} \mathrm{S} $$ If you begin with \(10.0 \mathrm{g}\) of \(\mathrm{CH}_{3} \mathrm{SH}\), and excess \(\mathrm{CO}\), (a) What is the theoretical yield of \(\mathrm{CH}_{3} \mathrm{COSCH}_{3} ?\) (b) If 8.65 g of \(\mathrm{CH}_{3} \mathrm{COSCH}_{3}\) is isolated, what is its percent yield?

A mixture of \(\mathrm{CuSO}_{4}\) and \(\mathrm{CuSO}_{4} \cdot 5 \mathrm{H}_{2} \mathrm{O}\) has a mass of 1.245 g. After heating to drive off all the water, the mass is only \(0.832 \mathrm{g} .\) What is the mass percent of \(\mathrm{CuSO}_{4} \cdot 5 \mathrm{H}_{2} \mathrm{O}\) in the mixture? (See page \(129 .\) )

A \(2.634-\mathrm{g}\) sample containing \(\mathrm{CuCl}_{2} \cdot 2 \mathrm{H}_{2} \mathrm{O}\) and other materials was heated. The sample mass after heating to drive off the water was \(2.125 \mathrm{g} .\) What was the mass percent of \(\mathrm{CuCl}_{2} \cdot 2 \mathrm{H}_{2} \mathrm{O}\) in the original sample?

A sample of limestone and other soil materials is heated, and the limestone decomposes to give calcium oxide and carbon dioxide. $$ \mathrm{CaCO}_{3}(\mathrm{s}) \longrightarrow \mathrm{CaO}(\mathrm{s})+\mathrm{CO}_{2}(\mathrm{g}) $$ A \(1.506-\mathrm{g}\) sample of limestone-containing material gives \(0.558 \mathrm{g}\) of \(\mathrm{CO}_{2},\) in addition to \(\mathrm{CaO},\) after being heated at a high temperature. What is the mass percent of \(\mathrm{CaCO}_{3}\) in the original sample?

At higher temperatures \(\mathrm{NaHCO}_{3}\) is converted quantitatively to \(\mathrm{Na}_{2} \mathrm{CO}_{3}\) \(2 \mathrm{NaHCO}_{3}(\mathrm{s}) \longrightarrow \mathrm{Na}_{2} \mathrm{CO}_{3}(\mathrm{s})+\mathrm{CO}_{2}(\mathrm{g})+\mathrm{H}_{2} \mathrm{O}(\mathrm{g})\) Heating a 1.7184 -g sample of impure NaHCO \(_{3}\) gives \(=0.196 \mathrm{g}\) of \(\mathrm{CO}_{2} .\) What was the mass percent of \(\mathrm{NaHCO}_{3}\) in the original 1.7184 -g sample?

A pesticide contains thallium( \(I\) ) sulfate, \(T I_{2} S O_{4} .\) Dissolving a 10.20 -g sample of impure pesticide in water and adding sodium iodide precipitates \(6.1964 \mathrm{g}\) of thallium (I) iodide, TII. $$ \mathrm{Tl}_{2} \mathrm{SO}_{4}(\mathrm{aq})+\mathrm{NaI}(\mathrm{aq}) \longrightarrow \mathrm{TII}(\mathrm{s})+\mathrm{Na}_{2} \mathrm{SO}_{4}(\mathrm{aq}) $$ What is the mass percent of \(\mathrm{TI}_{2} \mathrm{SO}_{4}\) in the original \(10.20-\mathrm{g}\) sample?

A The aluminum in a 0.764 -g sample of an unknown material was precipitated as aluminum hydroxide, Al(OH) \(_{3}\) which was then converted to \(\mathrm{Al}_{2} \mathrm{O}_{3}\) by heating strongly. If \(0.127 \mathrm{g}\) of \(\mathrm{Al}_{2} \mathrm{O}_{3}\) is obtained from the \(0.764 \mathrm{-g}\) sample, what is the mass percent of aluminum in the sample?

Styrene, the building block of polystyrene, consists of only \(\mathrm{C}\) and \(\mathrm{H}\). If \(0.438 \mathrm{g}\) of styrene is burned in oxygen and produces \(1.481 \mathrm{g}\) of \(\mathrm{CO}_{2}\) and \(0.303 \mathrm{g}\) of \(\mathrm{H}_{2} \mathrm{O},\) what is the empirical formula of styrene?

Mesitylene is a liquid hydrocarbon. Burning 0.115 g of the compound in oxygen gives \(0.379 \mathrm{g}\) of \(\mathrm{CO}_{2}\) and \(0.1035 \mathrm{g}\) of \(\mathrm{H}_{2} \mathrm{O} .\) What is the empirical formula of mesitylene?

Cyclopentane is a simple hydrocarbon. If 0.0956 g of the compound is burned in oxygen, \(0.300 \mathrm{g}\) of \(\mathrm{CO}_{2}\) and \(0.123 \mathrm{g}\) of \(\mathrm{H}_{2} \mathrm{O}\) are isolated. (a) What is the empirical formula of cyclopentane? (b) If a separate experiment gave \(70.1 \mathrm{g} / \mathrm{mol}\) as the molar mass of the compound, what is its molecular formula?

Azulene is a beautiful blue hydrocarbon. If 0.106 g of the compound is burned in oxygen, \(0.364 \mathrm{g}\) of \(\mathrm{CO}_{2}\) and \(0.0596 \mathrm{g}\) of \(\mathrm{H}_{2} \mathrm{O}\) are isolated. (a) What is the empirical formula of azulene? (b) If a separate experiment gave \(128.2 \mathrm{g} / \mathrm{mol}\) as the molar mass of the compound, what is its molecular formula?

An unknown compound has the formula \(\mathrm{C}_{x} \mathrm{H}_{1} \mathrm{O}_{2}\). You burn \(0.1523 \mathrm{g}\) of the compound and isolate \(0.3718 \mathrm{g}\) of \(\mathrm{CO}_{2}\) and \(0.1522 \mathrm{g}\) of \(\mathrm{H}_{2} \mathrm{O}\). What is the empirical formula of the compound? If the molar mass is \(72.1 \mathrm{g} / \mathrm{mol}\), what is the molecular formula? (See Exercise 4.9.)

Nickel forms a compound with carbon monoxide, \(\mathrm{Ni}_{x}(\mathrm{CO})_{r}\). To determine its formula, you carefully heat a 0.0973 -g sample in air to convert the nickel to \(0.0426 \mathrm{g}\) of NiO and the CO to \(0.100 \mathrm{g}\) of \(\mathrm{CO}_{2} .\) What is the empirical formula of \(\mathrm{Ni}_{x}(\mathrm{CO})_{y} ?\)

To find the formula of a compound composed of iron and carbon monoxide, \(\mathrm{Fe}_{x}(\mathrm{CO})_{y}\), the compound is burned in pure oxygen to give \(\mathrm{Fe}_{2} \mathrm{O}_{3}\) and \(\mathrm{CO}_{2} .\) If you burn \(1.959 \mathrm{g}\) of \(\mathrm{Fe}_{x}(\mathrm{CO})_{y}\) and obtain \(0.799 \mathrm{g}\) of \(\mathrm{Fe}_{2} \mathrm{O}_{3}\) and \(2.200 \mathrm{g}\) of \(\mathrm{CO}_{2}\) what is the empirical formula of \(\mathrm{Fe}_{x}(\mathrm{CO})_{y^{2}}\)

Balance the following equations: (a) The synthesis of urea, a common fertilizer \(\mathrm{CO}_{2}(\mathrm{g})+\mathrm{NH}_{3}(\mathrm{g}) \longrightarrow \mathrm{NH}_{2} \mathrm{CONH}_{2}(\mathrm{s})+\mathrm{H}_{2} \mathrm{O}(\ell)\) (b) Reactions used to make uranium(VI) fluoride for the enrichment of natural uranium \(\mathrm{UO}_{2}(\mathrm{s})+\mathrm{HF}(\mathrm{aq}) \longrightarrow \mathrm{UF}_{4}(\mathrm{s})+\mathrm{H}_{2} \mathrm{O}(\ell)\) \(\mathrm{UF}_{4}(\mathrm{s})+\mathrm{F}_{2}(\mathrm{g}) \longrightarrow \mathrm{UF}_{6}(\mathrm{s})\) (c) The reaction to make titanium(IV) chloride, which is then converted to titanium metal \(\mathrm{TiO}_{2}(\mathrm{s})+\mathrm{Cl}_{2}(\mathrm{g})+\mathrm{C}(\mathrm{s}) \longrightarrow \mathrm{TiCl}_{4}(\ell)+\mathrm{CO}(\mathrm{g})\) $$ \mathrm{TiCl}_{4}(\ell)+\mathrm{Mg}(\mathrm{s}) \longrightarrow \mathrm{Ti}(\mathrm{s})+\mathrm{MgCl}_{2}(\mathrm{s}) $$

Balance the following equations: (a) Reaction to produce "superphosphate" fertilizer \(\mathrm{Ca}_{3}\left(\mathrm{PO}_{4}\right)_{2}(\mathrm{s})+\mathrm{H}_{2} \mathrm{SO}_{4}(\mathrm{aq}) \longrightarrow \mathrm{Ca}\left(\mathrm{H}_{2} \mathrm{PO}_{4}\right)_{2}(\mathrm{aq})+\mathrm{CaSO}_{4}(\mathrm{s})\) (b) Reaction to produce diborane, \(B_{2} H_{6}\) \(\mathrm{NaBH}_{4}(\mathrm{s})+\mathrm{H}_{2} \mathrm{SO}_{4}(\mathrm{aq}) \longrightarrow \mathrm{B}_{2} \mathrm{H}_{6}(\mathrm{g})+\mathrm{H}_{2}(\mathrm{g})+\mathrm{Na}_{2} \mathrm{SO}_{4}(\mathrm{aq})\) (c) Reaction to produce tungsten metal from tungsten (VI) oxide \(\mathrm{WO}_{3}(\mathrm{s})+\mathrm{H}_{2}(\mathrm{g}) \longrightarrow \mathrm{W}(\mathrm{s})+\mathrm{H}_{2} \mathrm{O}(\ell)\) (d) Decomposition of ammonium dichromate \(\left(\mathrm{NH}_{4}\right)_{2} \mathrm{Cr}_{2} \mathrm{O}_{7}(\mathrm{s}) \longrightarrow \mathrm{N}_{2}(\mathrm{g})+\mathrm{H}_{2} \mathrm{O}(\ell)+\mathrm{Cr}_{2} \mathrm{O}_{3}(\mathrm{s})\)

Suppose \(16.04 \mathrm{g}\) of benzene, \(\mathrm{C}_{6} \mathrm{H}_{6},\) is burned in oxygen. (a) What are the products of the reaction? (b) What is the balanced equation for the reaction? (c) What mass of \(\mathrm{O}_{2}\), in grams, is required for complete combustion of benzene? (d) What is the total mass of products expected from \(16.04 \mathrm{g}\) of benzene?

If \(10.0 \mathrm{g}\) of carbon is combined with an exact, stoichiometric amount of oxygen \((26.6 \mathrm{g})\) to produce carbon dioxide, what is the theoretical yield of \(\mathrm{CO}_{2}\), in grams?

The metabolic disorder diabetes causes a buildup of acetone, \(\mathrm{CH}_{3} \mathrm{COCH}_{3}\), in the blood. Acetone, a volatile compound, is exhaled, giving the breath of untreated diabetics a distinctive odor. The acetone is produced by a breakdown of fats in a series of reactions. The equation for the last step is $$ \mathrm{CH}_{3} \mathrm{COCH}_{2} \mathrm{CO}_{2} \mathrm{H} \longrightarrow \mathrm{CH}_{3} \mathrm{COCH}_{3}+\mathrm{CO}_{2} $$ What mass of acetone can be produced from \(125 \mathrm{mg}\) of acetoacetic acid \(\left(\mathrm{CH}_{3} \mathrm{COCH}_{2} \mathrm{CO}_{2} \mathrm{H}\right) ?\)

Your body deals with excess nitrogen by excreting it in the form of urea, \(\mathrm{NH}_{2} \mathrm{CONH}_{2}\). The reaction producing it is the combination of arginine \(\left(\mathrm{C}_{6} \mathrm{H}_{14} \mathrm{N}_{4} \mathrm{O}_{2}\right)\) with water to give urea and ornithine \(\left(\mathrm{C}_{5} \mathrm{H}_{12} \mathrm{N}_{2} \mathrm{O}_{2}\right)\) \(\mathrm{C}_{6} \mathrm{H}_{14} \mathrm{N}_{4} \mathrm{O}_{2}+\mathrm{H}_{2} \mathrm{O} \longrightarrow \mathrm{NH}_{2} \mathrm{CONH}_{2}+\mathrm{C}_{5} \mathrm{H}_{12} \mathrm{N}_{2} \mathrm{O}_{2}\) Arginine Urea Ornithine If you excrete 95 mg of urea, what mass of arginine must have been used? What mass of ornithine must have been produced?

The reaction of \(750 .\) g each of \(\mathrm{NH}_{3}\) and \(\mathrm{O}_{2}\) was found to produce \(562 \mathrm{g}\) of \(\mathrm{NO}\) (see pages \(153-155\) ). $$ 4 \mathrm{NH}_{3}(\mathrm{g})+5 \mathrm{O}_{2}(\mathrm{g}) \longrightarrow 4 \mathrm{NO}(\mathrm{g})+6 \mathrm{H}_{2} \mathrm{O}(\mathrm{g}) $$ (a) What mass of water is produced by this reaction? (b) What quantity of \(\mathrm{O}_{2}\) is required to consume \(750 .\) g of \(\mathrm{NH}_{3} ?\)

Sodium azide, the explosive chemical used in automobile airbags, is made by the following reaction: $$ \mathrm{NaNO}_{3}+3 \mathrm{NaNH}_{2} \longrightarrow \mathrm{NaN}_{3}+3 \mathrm{NaOH}+\mathrm{NH}_{3} $$ If you combine \(15.0 \mathrm{g}\) of \(\mathrm{NaNO}_{3}(85.0 \mathrm{g} / \mathrm{mol})\) with \(15.0 \mathrm{g}\) of \(\mathrm{NaNH}_{2},\) what mass of \(\mathrm{NaN}_{3}\) is produced?

Iodine is made by the reaction $$\begin{aligned} 2 \mathrm{NaIO}_{3}(\mathrm{aq})+5 \mathrm{NaHSO}_{3}(\mathrm{aq}) & \longrightarrow \\\3 \mathrm{NaHSO}_{4}(\mathrm{aq})+2 \mathrm{Na}_{2} \mathrm{SO}_{4}(\mathrm{aq}) &+\mathrm{H}_{2} \mathrm{O}(\ell)+\mathrm{I}_{2}(\mathrm{aq}) \end{aligned}$$ (a) Name the two reactants. (b) If you wish to prepare \(1.00 \mathrm{kg}\) of \(\mathrm{I}_{2},\) what mass of NalO \(_{3}\) is required? What mass of \(\mathrm{NaHSO}_{3} ?\)

Copper(I) sulfide reacts with \(\mathrm{O}_{2}\) upon heating to give copper metal and sulfur dioxide. (a) Write a balanced equation for the reaction. (b) What mass of copper metal can be obtained from \(500 . \mathrm{g}\) of copper \((\mathrm{I})\) sulfide?