Chapter 7

In the combination reaction \(\mathrm{A}+2 \mathrm{B} \rightarrow \mathrm{C}, 1.00\) gram of substance \(A\) and 4.00 grams of substance \(B\) are consumed. How many grams of substance C are formed?

In the reaction \(A+B \rightarrow C+D, 3.00\) grams of substance \(C\) and 3.00 grams of substance \(\mathrm{D}\) are produced as 2.00 grams of substance A are consumed. How many grams of substance B are also consumed?

Depending on reaction conditions, \(\mathrm{O}_{2}\) may combine with \(\mathrm{N}_{2}\) to form \(\mathrm{NO}\) or \(\mathrm{NO}_{2} .\) If \(x\) grams of \(\mathrm{O}_{2}\) combine with \(y\) grams of \(\mathrm{N}_{2}\) to form \(\mathrm{NO}\), how many grams of \(\mathrm{O}_{2}\) combine with \(y\) grams of \(\mathrm{N}_{2}\) to form \(\mathrm{NO}_{2} ?\)

Combustion of sulfur may, depending on reaction conditions, produce \(\mathrm{SO}_{2}\) or \(\mathrm{SO}_{3}\). If \(x\) grams of \(\mathrm{O}_{2}\) combine with \(y\) grams of sulfur to form \(\mathrm{SO}_{2},\) how many grams of \(\mathrm{O}_{2}\) combine with \(y\) grams of sulfur to form \(\mathrm{SO}_{3} ?\)

Two of the more common oxides of iron have the formulas \(\mathrm{FeO}\) and \(\mathrm{Fe}_{2} \mathrm{O}_{3} .\) How much more oxygen combines with a given mass of iron to form \(\mathrm{Fe}_{2} \mathrm{O}_{3}\) than combines with the same mass of iron to form FeO?

Tin(IV) chloride is prepared by the following combination reaction: $$\mathrm{Sn}(s)+2 \mathrm{Cl}_{2}(g) \rightarrow \mathrm{SnCl}_{4}(\ell)$$ If \(x\) grams of chlorine combine with \(y\) grams of tin to form \(\operatorname{tin}(\mathrm{IV})\) chloride, how many grams of chlorine are there in a sample of \(\operatorname{tin}(\mathrm{II})\) chloride that contains \(y\) grams of tin?

In a balanced chemical equation, does the number of atoms in the reactants always equal the number of atoms in the products?

In a balanced chemical equation for the complete combustion of a hydrocarbon, what is the ratio of atoms of C in the hydrocarbon to molecules of \(\mathrm{CO}_{2}\) produced?

How many moles of water vapor are produced for every mole of methane consumed in the combustion reaction \(\mathrm{CH}_{4}(g)+2 \mathrm{O}_{2}(g) \rightarrow \mathrm{CO}_{2}(g)+2 \mathrm{H}_{2} \mathrm{O}(g) ?\)

Why is CO produced during the combustion of hydrocarbons even when enough \(\mathrm{O}_{2}\) is present for complete combustion?

Balance the following reactions for the formation of nitrogen compounds: a. \(\mathrm{N}_{2}(g)+\mathrm{O}_{2}(g) \rightarrow \mathrm{NO}(g)\) b. \(\mathrm{N}_{2}(g)+\mathrm{O}_{2}(g) \rightarrow \mathrm{N}_{2} \mathrm{O}(g)\) c. \(\mathrm{NO}(g)+\mathrm{NO}_{3}(g) \rightarrow \mathrm{NO}_{2}(g)\) d. \(\mathrm{NO}(g)+\mathrm{O}_{2}(g)+\mathrm{H}_{2} \mathrm{O}(\ell) \rightarrow \mathrm{HNO}_{2}(\ell)\)

Chemistry of Geothermal Vents Some scientists believe that life on Earth originated near geothermal vents. Balance the following reactions, which are among those taking place near such vents: a. \(\mathrm{CH}_{3} \mathrm{SH}(a q)+\mathrm{CO}(a q) \rightarrow \mathrm{CH}_{3} \mathrm{COSCH}_{3}(a q)+\mathrm{H}_{2} \mathrm{S}(a q)\) b. \(\mathrm{H}_{2} \mathrm{S}(a q)+\mathrm{CO}(a q) \rightarrow \mathrm{CH}_{3} \mathrm{CO}_{2} \mathrm{H}(a q)+\mathrm{S}_{8}(s)\)

Write a balanced chemical equation for each of the following reactions: a. Dinitrogen pentoxide reacts with sodium metal to produce sodium nitrate and nitrogen dioxide. b. A mixture of nitric acid and nitrous acid is formed when water reacts with dinitrogen tetroxide. c. At high pressure, nitrogen monoxide decomposes to dinitrogen monoxide and nitrogen dioxide.

Write a balanced chemical equation for each of the following reactions: a. Hydrazine, \(\mathrm{N}_{2} \mathrm{H}_{4},\) reacts with oxygen to produce water and the element nitrogen. b. Ammonia (NH_) burns in oxygen to produce elemental nitrogen and water. c. Silicon dioxide reacts with carbon to produce the element silicon and carbon monoxide.

Complete and balance the following chemical equations for the complete combustion of several hydrocarbons. a. \(\mathrm{C}_{3} \mathrm{H}_{8}(g)+\mathrm{O}_{2}(g) \rightarrow\) b. \(\mathrm{C}_{4} \mathrm{H}_{10}(g)+\mathrm{O}_{2}(g) \rightarrow\) c. \(C_{6} H_{6}(\ell)+O_{2}(g) \rightarrow\) d. \(\mathrm{C}_{8} \mathrm{H}_{18}(\ell)+\mathrm{O}_{2}(g) \rightarrow\)

Complete and balance the following chemical equations for the complete combustion of several hydrocarbons. a. \(\mathrm{C}_{5} \mathrm{H}_{10}(\ell)+\mathrm{O}_{2}(g) \rightarrow\) b. \(\mathrm{C}_{6} \mathrm{H}_{14}(\ell)+\mathrm{O}_{2}(g) \rightarrow\) c. \(\mathrm{C}_{8} \mathrm{H}_{10}(\ell)+\mathrm{O}_{2}(g) \rightarrow\) d. \(\mathrm{C}_{9} \mathrm{H}_{12}(\ell)+\mathrm{O}_{2}(g) \rightarrow\)

Chemistry of Volcanic Gases Balance the following reactions that occur during volcanic eruptions: a. \(\mathrm{SO}_{2}(g)+\mathrm{O}_{2}(g) \rightarrow \mathrm{SO}_{3}(g)\) b. \(\mathrm{H}_{2} \mathrm{S}(g)+\mathrm{O}_{2}(g) \rightarrow \mathrm{SO}_{2}(g)+\mathrm{H}_{2} \mathrm{O}(g)\) c. \(\mathrm{H}_{2} \mathrm{S}(g)+\mathrm{SO}_{2}(g) \rightarrow \mathrm{S}_{8}(s)+\mathrm{H}_{2} \mathrm{O}(g)\)

Copper was one of the first metals used by humans because it can be recovered from several copper minerals, including cuprite (Cu \(_{2} \mathrm{O}\) ), chalcocite (Cu \(_{2} \mathrm{S}\) ), and malachite \(\left[\mathrm{Cu}_{2} \mathrm{CO}_{3}(\mathrm{OH})_{2}\right] .\) Balance the following reactions for converting these minerals into copper metal: a. \(\mathrm{Cu}_{2} \mathrm{O}(s)+\mathrm{C}(s) \rightarrow \mathrm{Cu}(s)+\mathrm{CO}_{2}(g)\) b. \(\mathrm{Cu}_{2} \mathrm{O}(s)+\mathrm{Cu}_{2} \mathrm{S}(s) \rightarrow \mathrm{Cu}(s)+\mathrm{SO}_{2}(g)\) c. \(\mathrm{Cu}_{2} \mathrm{CO}_{3}(\mathrm{OH})_{2}(s)+\mathrm{C}(s) \rightarrow \mathrm{Cu}(s)+\mathrm{CO}_{2}(g)+\mathrm{H}_{2} \mathrm{O}(g)\)

There are two ways to write the equation for the combustion of ethane: $$\begin{aligned}\mathrm{C}_{2} \mathrm{H}_{6}(g)+\frac{7}{2} \mathrm{O}_{2}(g) \rightarrow 3 \mathrm{H}_{2} \mathrm{O}(g)+2 \mathrm{CO}_{2}(g) \\\2 \mathrm{C}_{2} \mathrm{H}_{6}(g)+7 \mathrm{O}_{2}(g) \rightarrow 6 \mathrm{H}_{2} \mathrm{O}(g)+4 \mathrm{CO}_{2}(g)\end{aligned}$$ Do the different ways of writing the equation affect the calculation of how much \(\mathrm{CO}_{2}\) is produced from a known quantity of \(\mathrm{C}_{2} \mathrm{H}_{6} ?\)

Suppose the same mass of each of these components of natural gas was completely combusted: (a) \(\mathrm{CH}_{4},\) (b) \(\mathrm{C}_{2} \mathrm{H}_{6}\) (c) \(\mathrm{C}_{3} \mathrm{H}_{8},\) and (d) \(\mathrm{C}_{4} \mathrm{H}_{10} .\) Which one would produce the greatest mass of \(\mathrm{CO}_{2} ?\)

Land Management It has been estimated that better management of cropland, grazing land, and forests could reduce the amount of carbon dioxide in the atmosphere by \(5.4 \times 10^{9}\) kilograms of carbon per year. a. How many moles of carbon are present in \(5.4 \times 10^{9}\) kilograms of carbon? b. How many kilograms of carbon dioxide does this quantity of carbon represent?

Most of the \(\mathrm{CO}_{2}\) emitted by industrial sources comes from the combustion of fossil fuels, but some is also produced by cement manufacturing and the conversion of limestone \(\left(\mathrm{CaCO}_{3}\right)\) to lime \((\mathrm{CaO}):\) $$\mathrm{CaCO}_{3}(s) \rightarrow \mathrm{CaO}(s)+\mathrm{CO}_{2}(g)$$ How many tons of \(\mathrm{CO}_{2}\) are produced per ton of lime?

When \(\mathrm{NaHCO}_{3}\) is heated above \(270^{\circ} \mathrm{C},\) it decomposes to \(\mathrm{Na}_{2} \mathrm{CO}_{3}(s), \mathrm{H}_{2} \mathrm{O}(g),\) and \(\mathrm{CO}_{2}(g)\) a. Write a balanced chemical equation for the decomposition reaction. b. Calculate the mass of \(\mathrm{CO}_{2}\) produced from the decomposition of \(25.0 \mathrm{g} \mathrm{NaHCO}_{3}.\)

Pigments for Stoplights Cadmium yellow (cadmium sulfide) is a lemon-yellow pigment used in the lenses of stoplights. Its formula is CdS, and it is very insoluble in water. The recommended recipe for cadmium yellow is to mix cadmium nitrate with sodium sulfide in water. The cadmium yellow forms as a solid, while the other product, sodium nitrate, remains dissolved in the water. a. Write a balanced chemical equation for the reaction. b. Calculate the mass of cadmium nitrate you must start with to make 125 g of CdS.

One step in the conversion of aluminum ore into aluminum metal involves the synthesis of cryolite \(\left(\mathrm{Na}_{3} \mathrm{AlF}_{6}\right)\) in the following reaction: $$6 \mathrm{HF}(g)+3 \mathrm{NaAlO}_{2}(s) \rightarrow \mathrm{Na}_{3} \mathrm{AlF}_{6}(s)+3 \mathrm{H}_{2} \mathrm{O}(\ell)+\mathrm{Al}_{2} \mathrm{O}_{3}(s)$$ How much NaAIO \(_{2}\) (sodium aluminate) is required to produce \(1.00 \mathrm{kg}\) of \(\mathrm{Na}_{3} \mathrm{AlF}_{6} ?\)

Chromium metal can be produced from high-temperature reactions of chromium(III) oxide with silicon or aluminum: $$\begin{aligned}\mathrm{Cr}_{2} \mathrm{O}_{3}(s)+2 \mathrm{Al}(\ell) & \rightarrow 2 \mathrm{Cr(\ell)+\mathrm{Al}_{2} \mathrm{O}_{3}(s) \\\2 \mathrm{Cr}_{2} \mathrm{O}_{3}(s)+3 \mathrm{Si}(\ell) & \rightarrow 4 \mathrm{Cr}(\ell)+3\mathrm{SiO}_{2}(s)\end{aligned}$$ a. Calculate the mass of aluminum required to prepare 400.0 grams of chromium metal by the first reaction. b. Calculate the mass of silicon required to prepare 400.0 grams of chromium metal by the second reaction.

Oxygen for First Responders In self-contained breathing devices used by first responders, potassium superoxide, \(\mathrm{KO}_{2}\) reacts with exhaled carbon dioxide to produce potassium carbonate and oxygen: $$4 \mathrm{KO}_{2}(s)+2 \mathrm{CO}_{2}(g) \rightarrow 2 \mathrm{K}_{2} \mathrm{CO}_{3}(s)+3 \mathrm{O}_{2}(g)$$ How much \(\mathrm{O}_{2}\) could be produced from \(85 \mathrm{g} \mathrm{KO}_{2} ?\)

Capturing \(\mathrm{CO}_{2}\) Carbon dioxide can be removed from a gas stream by reacting it with potassium carbonate in the presence of water: $$\mathrm{CO}_{2}(g)+\mathrm{K}_{2} \mathrm{CO}_{3}(s)+\mathrm{H}_{2} \mathrm{O}(\ell) \rightarrow 2 \mathrm{KHCO}_{3}(s)$$ If a resting human exhales \(36 \mathrm{mg} \mathrm{CO}_{2}\) in one breath, then how much potassium carbonate would be required to capture it all?

The uranium minerals found in nature must be refined and enriched in \(^{235} \mathrm{U}\) before the uranium can be used as a fuel in nuclear reactors. One procedure for enriching uranium relies on the reaction of UO, with HF at high temperatures to form UF \(_{4},\) which is then converted into UF \(_{6}\) in another high-temperature reaction with fluorine gas: $$(1)\mathrm{UO}_{2}(g)+4 \mathrm{HF}(a q) \rightarrow \mathrm{UF}_{4}(g)+2 \mathrm{H}_{2} \mathrm{O}(\ell)$$ $$(2)\mathrm{UF}_{4}(g)+\mathrm{F}_{2}(g) \rightarrow \mathrm{UF}_{6}(g)$$ a. How many kilograms of HF are needed to completely react with \(5.00 \mathrm{kg} \mathrm{UO}_{2} ?\) b. How much UF \(_{6}\) can be produced from \(850.0 \mathrm{g} \mathrm{UO}_{2} ?\)

Chalcopyrite (CuFeS \(_{2}\) ) is an abundant copper mineral that can be converted into elemental copper. How much Cu is there in \(1.00 \mathrm{kg} \mathrm{CuFeS}_{2} ?\)

Mining for Gold Unlike most metals, gold occurs in nature as the pure element. Miners in California in 1849 searched for gold nuggets and gold dust in streambeds, where the denser gold could be easily separated from sand and gravel. However, larger deposits of gold are found in veins of rock and can be separated chemically in the following two-step process: (1) \(4 \mathrm{Au}(s)+8 \mathrm{NaCN}(a q)+\mathrm{O}_{2}(g)+2 \mathrm{H}_{2} \mathrm{O}(\ell) \rightarrow\) $$4 \mathrm{NaAu}(\mathrm{CN})_{2}(a q)+4 \mathrm{NaOH}(a q)$$ (2) \(2 \mathrm{NaAu}(\mathrm{CN})_{2}(a q)+\mathrm{Zn}(s) \rightarrow\) $$2 \mathrm{Au}(s)+\mathrm{Na}_{2}\left[\mathrm{Zn}(\mathrm{CN})_{4}\right](a q)$$ If 23 kilograms of ore is \(0.19 \%\) gold by mass, how much Zn is needed to react with the gold in the ore? Assume that reactions 1 and 2 are \(100 \%\) efficient.

What is the difference between an empirical formula and a molecular formula?

Do the empirical and molecular formulas of a compound have the same percent composition values?

Sometimes the composition of a compound is expressed in units of mol \(\%\) or atom \(\% .\) Are the values of these parameters likely to be the same or different for a given compound?

Calculate the percent composition of (a) \(\mathrm{Na}_{2} \mathrm{O},\) (b) \(\mathrm{NaOH}\) (c) \(\mathrm{NaHCO}_{3},\) and \((\mathrm{d}) \mathrm{Na}_{2} \mathrm{CO}_{3}.\)

Calculate the percent composition of (a) sodium sulfate, (b) dinitrogen tetroxide, (c) strontium nitrate, and (d) aluminum sulfide.

Organic Compounds in Space The following compounds have been detected in space. Which contains the greatest percentage of carbon by mass? a. naphthalene, \(\mathrm{C}_{10} \mathrm{H}_{8}\) b. chrysene, \(\mathrm{C}_{18} \mathrm{H}_{12}\) c. pentacene, \(\mathrm{C}_{22} \mathrm{H}_{14}\) d. pyrene, \(C_{16} H_{10}\)

Lead Compounds as Pigments Ancient Egyptians used lead compounds, including \(\mathrm{PbS}, \mathrm{PbCO}_{3},\) and \(\mathrm{Pb}_{2} \mathrm{Cl}_{2} \mathrm{CO}_{3}\) as pigments in cosmetics, and many people suffered from chronic lead poisoning as a result. Calculate the percentage of lead in each of the compounds.

A 3.556 g sample of a pure aluminum oxide decomposes under high heat to produce \(1.674 \mathrm{g}\) of oxygen in addition to pure aluminum metal. What is the empirical formula of the aluminum oxide?

Oxygen gas may be prepared in the laboratory by decomposing a compound of potassium, chlorine, and oxygen. The other product of the decomposition is \(\mathrm{KCl}(s)\) Complete decomposition of \(2.917 \mathrm{g}\) of the compound produces \(1.143 \mathrm{g}\) of oxygen gas. What is the empirical formula of the compound?

Do any two of the following compounds, which have been detected in outer space, have the same empirical formula? a. naphthalene, \(\mathrm{C}_{10} \mathrm{H}_{8}\) b. chrysene, \(\mathbf{C}_{18} \mathrm{H}_{12}\) c. anthracene, \(\mathrm{C}_{14} \mathrm{H}_{10}\) d. pyrene, \(\mathrm{C}_{16} \mathrm{H}_{10}\) e. benzoperylene, \(\mathrm{C}_{22} \mathrm{H}_{12}\) f. coronene, \(\mathrm{C}_{24} \mathrm{H}_{12}\)

Which of the following nitrogen oxides have the same empirical formulas? (a) \(\mathrm{N}_{2} \mathrm{O} ;\) (b) \(\mathrm{NO} ;\) (c) \(\mathrm{NO}_{2} ;\) (d) \(\mathrm{N}_{2} \mathrm{O}_{2}\) (e) \(\mathrm{N}_{2} \mathrm{O}_{4}\)

Surgical-Grade Titanium Medical implants and high-quality jewelry items for body piercings are frequently made of a material known as G23Ti, or surgical- grade titanium. The percent composition of the material is \(64.39 \%\) titanium, \(24.19 \%\) aluminum, and \(11.42 \%\) vanadium. What is the empirical formula of surgical-grade titanium?

A sample of an iron-containing compound is \(22.0 \%\) iron, \(50.2 \%\) oxygen, and \(27.8 \%\) chlorine by mass. What is the empirical formula of this compound?

Phosphorus burns in pure oxygen with a brilliant white light. The product of combustion is \(43.64 \%\) phosphorus and \(56.36 \%\) oxygen. a. What is the empirical formula of this compound? b. The molar mass of the compound is \(284 \mathrm{g} / \mathrm{mol}\). What is its molecular formula?

Chemistry of Soot A piece of glass held over a candle flame becomes coated with soot, which is the result of the incomplete combustion of candle wax. Elemental analysis of a compound extracted from a sample of soot gave the following results: \(7.74 \%\) H and \(92.26 \%\) C by mass. Calculate the empirical formula of the compound.

What is the empirical formula of the compound that is \(24.2 \% \mathrm{Cu}, 27.0 \% \mathrm{Cl},\) and \(48.8 \%\) O by mass?

A chlorine oxide used to kill anthrax spores in contaminated buildings is \(52.6 \%\) Cl by mass. What is its empirical formula?

Explain why it is important for combustion analysis to be carried out in an excess of oxygen.

Why is the quantity of \(\mathrm{CO}_{2}\) obtained in a combustion analysis not a direct measure of the oxygen content of the starting compound?