Chapter 4

Identify the reactants and products in this chemical equation. \(4 \mathrm{NH}_{3}(g)+5 \mathrm{O}_{2}(g) \longrightarrow 4 \mathrm{NO}(g)+6 \mathrm{H}_{2} \mathrm{O}(g)\).

Why must chemical equations be balanced?

What is reaction stoichiometry? What is the significance of the coefficients in a balanced chemical equation?

What is a combustion reaction? Why are combustion reactions important? Give an example.

Write a general equation for the reaction of an alkali metal with: a. a halogen b. water

Write a general equation for the reaction of a halogen with: a. a metal b. hydrogen c. another halogen

Sulfuric acid is a component of acid rain formed when gaseous sulfur dioxide pollutant reacts with gaseous oxygen and liquid water to form aqueous sulfuric acid. Write the balanced chemical equation for this reaction. (Note: This is a simplified representation of this reaction.)

Nitric acid is a component of acid rain that forms when gaseous nitrogen dioxide pollutant reacts with gaseous oxygen and liquid water to form aqueous nitric acid. Write the balanced chemical equation for this reaction. (Note: This is a simplified representation of this reaction.)

In a popular classroom demonstration, solid sodium is added to liquid water and reacts to produce hydrogen gas and aqueous sodium hydroxide. Write the balanced chemical equation for this reaction.

When iron rusts, solid iron reacts with gaseous oxygen to form solid iron(III) oxide. Write the balanced chemical equation for this reaction.

Write the balanced chemical equation for the fermentation of sucrose \(\left(\mathrm{C}_{12} \mathrm{H}_{22} \mathrm{O}_{11}\right)\) by yeasts in which the aqueous sugar reacts with water to form aqueous ethanol \(\left(\mathrm{C}_{2} \mathrm{H}_{5} \mathrm{OH}\right)\) and carbon dioxide gas.

Write the balanced chemical equation for each reaction. a. Solid lead(II) sulfide reacts with aqueous hydrobromic acid to form solid lead(II) bromide and dihydrogen monosulfide gas. b. Gaseous carbon monoxide reacts with hydrogen gas to form gaseous methane \(\left(\mathrm{CH}_{4}\right)\) and liquid water. c. Aqueous hydrochloric acid reacts with solid manganese(IV) oxide to form aqueous manganese(II) chloride, liquid water, and chlorine gas. d. Liquid pentane \(\left(\mathrm{C}_{5} \mathrm{H}_{12}\right)\) reacts with gaseous oxygen to form carbon dioxide and liquid water.

Write the balanced chemical equation for each reaction. a. Solid copper reacts with solid sulfur to form solid copper(I) sulfide. b. Solid iron(III) oxide reacts with hydrogen gas to form solid iron and liquid water. c. Sulfur dioxide gas reacts with oxygen gas to form sulfur trioxide gas. d. Gaseous ammonia \(\left(\mathrm{NH}_{3}\right)\) reacts with gaseous oxygen to form gaseous nitrogen monoxide and gaseous water.

Write the balanced chemical equation for the reaction of aqueous sodium carbonate with aqueous copper(II) chloride to form solid copper(II) carbonate and aqueous sodium chloride.

Write the balanced chemical equation for the reaction of aqueous potassium hydroxide with aqueous iron(III) chloride to form solid iron(III) hydroxide and aqueous potassium chloride.

Balance each chemical equation. a. \(\mathrm{CO}_{2}(g)+\mathrm{CaSiO}_{3}(s)+\mathrm{H}_{2} \mathrm{O}(l) \longrightarrow \mathrm{SiO}_{2}(s)+\mathrm{Ca}\left(\mathrm{HCO}_{3}\right)_{2}(a q)\) b. \(\mathrm{Co}\left(\mathrm{NO}_{3}\right)_{3}(a q)+\left(\mathrm{NH}_{4}\right)_{2} \mathrm{~S}(a q) \longrightarrow \mathrm{Co}_{2} \mathrm{~S}_{3}(s)+\mathrm{NH}_{4} \mathrm{NO}_{3}(a q)\) c. \(\mathrm{Cu}_{2} \mathrm{O}(s)+\mathrm{C}(s) \longrightarrow \mathrm{Cu}(s)+\mathrm{CO}(g)\) d. \(\mathrm{H}_{2}(g)+\mathrm{Cl}_{2}(g) \longrightarrow \mathrm{HCl}(g)\)

Balance each chemical equation. a. \(\mathrm{Na}_{2} \mathrm{~S}(a q)+\mathrm{Cu}\left(\mathrm{NO}_{3}\right)_{2}(a q) \longrightarrow \mathrm{NaNO}_{3}(a q)+\mathrm{CuS}(s)\) b. \(\mathrm{N}_{2} \mathrm{H}_{4}(l) \longrightarrow \mathrm{NH}_{3}(g)+\mathrm{N}_{2}(g)\) c. \(\mathrm{HCl}(a q)+\mathrm{O}_{2}(g) \longrightarrow \mathrm{H}_{2} \mathrm{O}(l)+\mathrm{Cl}_{2}(g)\) d. \(\mathrm{FeS}(s)+\mathrm{HCl}(a q) \longrightarrow \mathrm{FeCl}_{2}(a q)+\mathrm{H}_{2} \mathrm{~S}(g)\)

Consider the unbalanced equation for the combustion of hexane: $$ \mathrm{C}_{6} \mathrm{H}_{14}(g)+\mathrm{O}_{2}(g) \longrightarrow \mathrm{CO}_{2}(g)+\mathrm{H}_{2} \mathrm{O}(g) $$ Balance the equation and determine how many moles of \(\mathrm{O}_{2}\) are required to react completely with 7.2 moles of \(\mathrm{C}_{6} \mathrm{H}_{14}\).

Consider the balanced equation: $$ \mathrm{SiO}_{2}(s)+3 \mathrm{C}(s) \longrightarrow \mathrm{SiC}(s)+2 \mathrm{CO}(g) $$ Complete the table showing the appropriate number of moles of reactants and products. If the number of moles of a reactant is provided, fill in the required amount of the other reactant, as well as the moles of each product that forms. If the number of moles of a product is provided, fill in the required amount of each reactant to make that amount of product, as well as the amount of the other product that forms.

Hydrobromic acid dissolves solid iron according to the reaction: $$ \mathrm{Fe}(s)+2 \mathrm{HBr}(a q) \longrightarrow \mathrm{FeBr}_{2}(a q)+\mathrm{H}_{2}(g) $$ What mass of HBr (in g) do you need to dissolve a 3.2 -g pure iron bar on a padlock? What mass of \(\mathrm{H}_{2}\) would the complete reaction of the iron bar produce?

Sulfuric acid dissolves aluminum metal according to the reaction: $$ 2 \mathrm{Al}(s)+3 \mathrm{H}_{2} \mathrm{SO}_{4}(a q) \longrightarrow \mathrm{Al}_{2}\left(\mathrm{SO}_{4}\right)_{3}(a q)+3 \mathrm{H}_{2}(g) $$ Suppose you want to dissolve an aluminum block with a mass of 15.2 g. What minimum mass of \(\mathrm{H}_{2} \mathrm{SO}_{4}\) (in g) do you need? What mass of \(\mathrm{H}_{2}\) gas (in g) does the complete reaction of the aluminum block produce?

Find the limiting reactant for each initial amount of reactants. $$ 2 \mathrm{Na}(s)+\mathrm{Br}_{2}(g) \longrightarrow 2 \mathrm{NaBr}(s) $$ a. \(2 \mathrm{~mol} \mathrm{Na}, 2 \mathrm{~mol} \mathrm{Br}_{2}\) b. \(1.8 \mathrm{~mol} \mathrm{Na}, 1.4 \mathrm{~mol} \mathrm{Br}_{2}\) c. \(2.5 \mathrm{~mol} \mathrm{Na}, 1 \mathrm{~mol} \mathrm{Br}_{2}\) d. \(12.6 \mathrm{~mol} \mathrm{Na}, 6.9 \mathrm{~mol} \mathrm{Br}_{2}\)

Find the limiting reactant for each initial amount of reactants. $$ 4 \mathrm{Al}(s)+3 \mathrm{O}_{2}(g) \longrightarrow 2 \mathrm{Al}_{2} \mathrm{O}_{3}(s) $$ a. \(1 \mathrm{~mol} \mathrm{Al}, 1 \mathrm{~mol} \mathrm{O}_{2}\) b. \(4 \mathrm{~mol} \mathrm{Al}, 2.6 \mathrm{~mol} \mathrm{O}_{2}\) c. \(16 \mathrm{~mol} \mathrm{Al}, 13 \mathrm{~mol} \mathrm{O}_{2}\) d. \(7.4 \mathrm{~mol} \mathrm{Al}, 6.5 \mathrm{~mol} \mathrm{O}_{2}\)

Calculate the theoretical yield of the product (in moles) for each initial amount of reactants. $$ \mathrm{Ti}(s)+2 \mathrm{Cl}_{2}(g) \longrightarrow \mathrm{TiCl}_{4}(l) $$ a. \(4 \mathrm{~mol} \mathrm{Ti}, 4 \mathrm{~mol} \mathrm{Cl}_{2}\) b. \(7 \mathrm{~mol} \mathrm{Ti}, 17 \mathrm{~mol} \mathrm{Cl}_{2}\) c. \(12.4 \mathrm{~mol} \mathrm{Ti}, 18.8 \mathrm{~mol} \mathrm{Cl}_{2}\)

Calculate the theoretical yield of product (in moles) for each initial amount of reactants. $$ 3 \mathrm{Mn}(s)+2 \mathrm{O}_{2}(g) \longrightarrow \mathrm{Mn}_{3} \mathrm{O}_{4}(s) $$ a. \(3 \mathrm{~mol} \mathrm{Mn}, 3 \mathrm{~mol} \mathrm{O}_{2}\) b. \(4 \mathrm{~mol} \mathrm{Mn}, 7 \mathrm{~mol} \mathrm{O}_{2}\) c. \(27.5 \mathrm{~mol} \mathrm{Mn}, 43.8 \mathrm{~mol} \mathrm{O}_{2}\)

Zinc sulfide reacts with oxygen according to the reaction: $$ 2 \mathrm{ZnS}(s)+3 \mathrm{O}_{2}(g) \longrightarrow 2 \mathrm{ZnO}(s)+2 \mathrm{SO}_{2}(g) $$ A reaction mixture initially contains \(4.2 \mathrm{~mol} \mathrm{ZnS}\) and \(6.8 \mathrm{~mol}\) \(\mathrm{O}_{2} .\) Once the reaction has occurred as completely as possible, what amount (in moles) of the excess reactant remains?

Iron(II) sulfide reacts with hydrochloric acid according to the reaction: $$ \mathrm{FeS}(s)+2 \mathrm{HCl}(a q) \longrightarrow \mathrm{FeCl}_{2}(s)+\mathrm{H}_{2} \mathrm{~S}(g) $$ A reaction mixture initially contains 0.223 mol FeS and 0.652 mol HCl. Once the reaction has occurred as completely as possible, what amount (in moles) of the excess reactant remains?

For the reaction shown, calculate the theoretical yield of product (in grams) for each initial amount of reactants. $$ 2 \mathrm{Al}(s)+3 \mathrm{Cl}_{2}(g) \longrightarrow 2 \mathrm{AlCl}_{3}(s) $$ a. \(2.0 \mathrm{~g} \mathrm{Al}, 2.0 \mathrm{~g} \mathrm{Cl}_{2}\) b. \(7.5 \mathrm{~g} \mathrm{Al}, 24.8 \mathrm{~g} \mathrm{Cl}_{2}\) c. \(0.235 \mathrm{~g} \mathrm{Al}, 1.15 \mathrm{~g} \mathrm{Cl}_{2}\)

Iron(III) oxide reacts with carbon monoxide according to the equation: $$ \mathrm{Fe}_{2} \mathrm{O}_{3}(s)+3 \mathrm{CO}(g) \longrightarrow 2 \mathrm{Fe}(s)+3 \mathrm{CO}_{2}(g) $$ A reaction mixture initially contains \(22.55 \mathrm{~g} \mathrm{Fe}_{2} \mathrm{O}_{3}\) and \(14.78 \mathrm{~g}\) CO. Once the reaction has occurred as completely as possible, what mass (in g) of the excess reactant remains?

Elemental phosphorus reacts with chlorine gas according to the equation: $$ \mathrm{P}_{4}(s)+6 \mathrm{Cl}_{2}(g) \longrightarrow 4 \mathrm{PCl}_{3}(l) $$ A reaction mixture initially contains \(45.69 \mathrm{~g} \mathrm{P}_{4}\) and \(131.3 \mathrm{~g} \mathrm{Cl}_{2}\). Once the reaction has occurred as completely as possible, what mass (in g) of the excess reactant remains?

Complete and balance each combustion reaction equation: a. \(\mathrm{C}_{4} \mathrm{H}_{6}(g)+\mathrm{O}_{2}(g) \longrightarrow\) b. \(\mathrm{C}(s)+\mathrm{O}_{2}(g) \longrightarrow\) c. \(\mathrm{CS}_{2}(s)+\mathrm{O}_{2}(g) \longrightarrow\) d. \(\mathrm{C}_{3} \mathrm{H}_{8} \mathrm{O}(l)+\mathrm{O}_{2}(g) \longrightarrow\)

Write a balanced chemical equation for the reaction of solid strontium with iodine gas.

Write a balanced chemical equation for the reaction between lithium metal and chlorine gas.

Write a balanced chemical equation for the reaction of solid lithium with liquid water.

Write a balanced chemical equation for the reaction of solid potassium with liquid water.

Write a balanced equation for the reaction of hydrogen gas with bromine gas.

Write a balanced equation for the reaction of chlorine gas with fluorine gas.

The combustion of gasoline produces carbon dioxide and water. Assume gasoline to be pure octane \(\left(\mathrm{C}_{8} \mathrm{H}_{18}\right)\) and calculate the mass (in kg) of carbon dioxide that is added to the atmosphere per \(1.0 \mathrm{~kg}\) of octane burned. (Hint: Begin by writing a balanced equation for the combustion reaction.

Aspirin can be made in the laboratory by reacting acetic anhydride \(\left(\mathrm{C}_{4} \mathrm{H}_{6} \mathrm{O}_{3}\right)\) with salicylic acid \(\left(\mathrm{C}_{7} \mathrm{H}_{6} \mathrm{O}_{3}\right)\) to form aspirin \(\left(\mathrm{C}_{9} \mathrm{H}_{8} \mathrm{O}_{4}\right)\) and acetic acid \(\left(\mathrm{C}_{2} \mathrm{H}_{4} \mathrm{O}_{2}\right) .\) The balanced equation is: $$ \mathrm{C}_{4} \mathrm{H}_{6} \mathrm{O}_{3}+\mathrm{C}_{7} \mathrm{H}_{6} \mathrm{O}_{3} \longrightarrow \mathrm{C}_{9} \mathrm{H}_{8} \mathrm{O}_{4}+\mathrm{C}_{2} \mathrm{H}_{4} \mathrm{O}_{2} $$ In a laboratory synthesis, a student begins with \(3.00 \mathrm{~mL}\) of acetic anhydride (density \(=1.08 \mathrm{~g} / \mathrm{mL}\) ) and \(1.25 \mathrm{~g}\) of salicylic acid. Once the reaction is complete, the student collects \(1.22 \mathrm{~g}\) of aspirin. Determine the limiting reactant, theoretical yield of aspirin, and percent yield for the reaction.

A mixture of A and B contains a total of 5.3 mols. Both A and B react with Z according to the following equations: $$ \begin{array}{c} \mathrm{A}+\mathrm{Z} \longrightarrow \mathrm{AZ} \\ \mathrm{B}+2 \mathrm{Z} \longrightarrow \mathrm{BZ}_{2} \end{array} $$ The reaction of the mixture of A and B with Z consumes \(7.8 \mathrm{~mol} \mathrm{Z}\). Assuming the reactions go to completion, how many moles of A does the mixture contain?

Consider the reaction: $$ 4 \mathrm{~K}(s)+\mathrm{O}_{2}(g) \longrightarrow 2 \mathrm{~K}_{2} \mathrm{O}(s) $$ The molar mass of \(\mathrm{K}\) is \(39.10 \mathrm{~g} / \mathrm{mol}\), and that of \(\mathrm{O}_{2}\) is \(32.00 \mathrm{~g} / \mathrm{mol}\). Without doing any calculations, pick the conditions under which potassium is the limiting reactant and explain your reasoning. a. \(170 \mathrm{~g} \mathrm{~K}, 31 \mathrm{~g} \mathrm{O}_{2}\) b. \(16 \mathrm{~g} \mathrm{~K}, 2.5 \mathrm{~g} \mathrm{O}_{2}\) c. \(165 \mathrm{~kg} \mathrm{~K}, 28 \mathrm{~kg} \mathrm{O}_{2}\) d. \(1.5 \mathrm{~g} \mathrm{~K}, 0.38 \mathrm{~g} \mathrm{O}_{2}\)

Consider the reaction: $$ 2 \mathrm{NO}(g)+5 \mathrm{H}_{2}(g) \longrightarrow 2 \mathrm{NH}_{3}(g)+2 \mathrm{H}_{2} \mathrm{O}(g) $$ A reaction mixture initially contains 5 moles of NO and 10 moles of \(\mathrm{H}_{2}\). Without doing any calculations, determine which set of amounts best represents the mixture after the reactants have reacted as completely as possible. Explain your reasoning. a. \(1 \mathrm{~mol} \mathrm{NO}, 0 \mathrm{~mol} \mathrm{H}_{2}, 4 \mathrm{~mol} \mathrm{NH}_{3}, 4 \mathrm{~mol} \mathrm{H}_{2} \mathrm{O}\) b. \(0 \mathrm{~mol} \mathrm{NO}, 1 \mathrm{~mol} \mathrm{H}_{2}, 5 \mathrm{~mol} \mathrm{NH}_{3}, 5 \mathrm{~mol} \mathrm{H}_{2} \mathrm{O}\) c. \(3 \mathrm{~mol} \mathrm{NO}, 5 \mathrm{~mol} \mathrm{H}_{2}, 2 \mathrm{~mol} \mathrm{NH}_{3}, 2 \mathrm{~mol} \mathrm{H}_{2} \mathrm{O}\) d. \(0 \mathrm{~mol} \mathrm{NO}, 0 \mathrm{~mol} \mathrm{H}_{2}, 4 \mathrm{~mol} \mathrm{NH}_{3}, 4 \mathrm{~mol} \mathrm{H}_{2} \mathrm{O}\)

Imagine you mix \(16.05 \mathrm{~g}\) of methane \(\left(\mathrm{CH}_{4}\right)\) gas and \(96.00 \mathrm{~g}\) of oxygen \(\left(\mathrm{O}_{2}\right)\) gas and then ignite the mixture. After a bright flash and a loud bang, some water vapor forms. a. Write the balanced chemical reaction for the combustion of methane. b. Depict the process that occurred using circles to represent atoms. Represent carbon with black circles, hydrogen with white circles, and oxygen with gray circles. Let one circle (or one molecule made of circles bonded together) represent exactly one mole. c. How many moles of water can you make? How many moles of carbon dioxide? d. Will anything be left over? If so, how much? e. Identify the following: limiting reagent, excess reagent, and theoretical yield.