Chapter 19

Explain why writing hydrogen ions as H + in redox reactions represents a simplification and not how they exist

Before you attempt to balance the equation for a redox reaction, why do you need to know whether the reaction takes place in acidic or basic solution?

Explain what a spectator ion is.

Define the term species in terms of redox reactions.

Does the following equation represent a reduction or an oxidation process? Explain your answer. $$\mathrm{Zn}^{2+}+2 \mathrm{e}^{-} \rightarrow \mathrm{Zn}$$

Describe what is happening to electrons in each half reaction of a redox process

Use the oxidation-number method to balance these redox equations. a. \(\mathrm{Cl}_{2}+\mathrm{NaOH} \rightarrow \mathrm{NaCl}+\mathrm{HOCl}\) b. \(\mathrm{HBrO}_{3} \rightarrow \mathrm{Br}_{2}+\mathrm{H}_{2} \mathrm{O}+\mathrm{O}_{2}\)

Balance these net ionic equations for redox reactions. a. \(A u^{3+}(a q)+I^{-(a q)} \rightarrow \operatorname{Au}(s)+I_{2}(s)\) b. \(C e^{4+}(a q)+S n^{2+}(a q) \rightarrow C e^{3+}(a q)+S n^{4+}(a q)\)

Use the oxidation-number method to balance the following ionic redox equations. a. \(A l+I_{2} \rightarrow A l^{3+}+I^{-}\) b. \(M n O_{2}+B r^{-} \rightarrow M n^{2+}+B r_{2}(\text { in acid solution })\)

Use the oxidation-number method to balance these redox equations. a. \(\mathrm{PbS}+\mathrm{O}_{2} \rightarrow \mathrm{PbO}+\mathrm{SO}_{2}\) b. NaWO \(_{3}+\mathrm{NaOH}+\mathrm{O}_{2} \rightarrow \mathrm{Na}_{2} \mathrm{WO}_{4}+\mathrm{H}_{2} \mathrm{O}\) c. \(\mathrm{NH}_{3}+\mathrm{CuO} \rightarrow \mathrm{Cu}+\mathrm{N}_{2}+\mathrm{H}_{2} \mathrm{O}\) d. \(\mathrm{Al}_{2} \mathrm{O}_{3}+\mathrm{C}+\mathrm{Cl}_{2} \rightarrow \mathrm{AlCl}_{3}+\mathrm{CO}\)

Write the oxidation and reduction half-reactions represented in each of these redox equations. Write the halfreactions in net ionic form if they occur in aqueous solution. a. \(\mathrm{PbO}(\mathrm{s})+\mathrm{NH}_{3}(\mathrm{g}) \rightarrow \mathrm{N}_{2}(\mathrm{g})+\mathrm{H}_{2} \mathrm{O}(\mathrm{l})+\mathrm{Pb}(\mathrm{s})\) b. \(\mathrm{I}_{2}(\mathrm{s})+\mathrm{Na}_{2} \mathrm{S}_{2} \mathrm{O}_{3}(\mathrm{aq}) \rightarrow \mathrm{Na}_{2} \mathrm{S}_{2} \mathrm{O}_{4}(\mathrm{aq})+\mathrm{NaI}(\mathrm{aq})\) c. \(\mathrm{Sn}(\mathrm{s})+2 \mathrm{HCl}(\mathrm{aq}) \rightarrow \mathrm{SnCl}_{2}(\mathrm{aq})+\mathrm{H}_{2}(\mathrm{g})\)

Write the two half-reactions that make up the following balanced redox reaction. \(3 \mathrm{H}_{2} \mathrm{C}_{2} \mathrm{O}_{4}+2 \mathrm{HAsO}_{2} \rightarrow 6 \mathrm{CO}_{2}+2 \mathrm{As}+4 \mathrm{H}_{2} \mathrm{O}\)

Label each half-reaction as reduction or oxidation. a. \(\mathrm{Fe}^{2+}(\mathrm{aq}) \rightarrow \mathrm{Fe}^{3+}(\mathrm{aq})+\mathrm{e}^{-}\) b. \(\mathrm{MnO}_{4}-5 \mathrm{e}^{-}+8 \mathrm{H}^{+} \rightarrow \mathrm{Mn}^{2+}+4 \mathrm{H}_{2} \mathrm{O}\) c. \(2 \mathrm{H}^{+}+2 \mathrm{e}^{-} \rightarrow \mathrm{H}_{2}\) d. \(\mathrm{F}_{2} \rightarrow 2 \mathrm{F}^{-}+2 \mathrm{e}^{-}\)

Copper When solid copper pieces are put into a solution of silver nitrate, as shown in Figure 19.12, silver metal appears and blue copper(II) nitrate forms. Write the corresponding chemical equation without balancing it. Next, determine the oxidation state of each element in the equation. Write the two half-reactions, labeling which is oxidation and which is reduction. Finally, write a balanced equation for the reaction.

Use the oxidation-number method to balance these ionic redox equations. a. \(\mathrm{MoCl}_{5}+\mathrm{S}^{2-} \rightarrow \mathrm{MoS}_{2}+\mathrm{Cl}^{-}+\mathrm{S}\) b. \(\mathrm{TiCl}_{6}^{2-}+\mathrm{Zn} \rightarrow \mathrm{Ti}^{3+}+\mathrm{Cl}^{-}+\mathrm{Zn}^{2+}\)

Use the half-reaction method to balance these equations for redox reactions. Add water molecules and hydrogen ions (in acid solutions) or hydroxide ions (in basic solutions) as needed a. \(\mathrm{NH}_{3}(\mathrm{g})+\mathrm{NO}_{2}(\mathrm{g}) \rightarrow \mathrm{N}_{2}(\mathrm{g})+\mathrm{H}_{2} \mathrm{O}(1)\) b. \(\mathrm{Br}_{2} \rightarrow \mathrm{Br}^{-}+\mathrm{BrO}_{3}^{-}(\text { in basic solution })\)

Balance the following redox chemical equation. Rewrite the equation in full ionic form, then derive the net ionic equation and balance by the half- reaction method. Give the final answer as it is shown below but with the balancing coefficients. \(\mathrm{KMnO}_{4}(\mathrm{aq})+\mathrm{FeSO}_{4}(\mathrm{aq})+\mathrm{H}_{2} \mathrm{SO}_{4}(\mathrm{aq}) \rightarrow\) \(\mathrm{Fe}_{2}\left(\mathrm{SO}_{4}\right)_{3}(\mathrm{aq})+\mathrm{MnSO}_{4}(\mathrm{aq})+\) \(\quad \mathrm{K}_{2} \mathrm{SO}_{4}(\mathrm{aq})+\mathrm{H}_{2} \mathrm{O}(1)\)

Write the oxidation and reduction half-reaction represented in each of these redox equations. Write the half-reactions in net ionic form if they occur in aqueous solution. a. \(P b O(s)+N H_{3}(g) \rightarrow N_{2}(g)+H_{2} O(1)+P b(s)\) b. \(I_{2}(s)+N a S_{2} O_{3}(a q) \rightarrow N a_{2} S_{2} O_{4}(a q)+N a I(a q)\) c. \(\operatorname{Sn}(s)+2 H C l(a q) \rightarrow \operatorname{Sn} C l_{2}(a q)+H_{2}(g)\)

Use the half-reaction method to balance these equations. Add water molecules and hydrogen ions (in acid solutions) or hydroxide ions (in basic solutions) as needed. Keep balanced equations in net ionic form. a. \(\mathrm{Cl}^{-}(\mathrm{aq})+\mathrm{NO}_{3}-(\mathrm{aq}) \rightarrow \mathrm{ClO}^{-}(\mathrm{aq})+\mathrm{NO}(\mathrm{g})\) (in acid solution) b. \(\mathrm{IO}_{3}-(\mathrm{aq})+\mathrm{Br}^{-}(\mathrm{aq}) \rightarrow \mathrm{Br}_{2}(\mathrm{l})+\mathrm{IBr}(\mathrm{s})\) (in acid solution) c. \(\mathrm{I}_{2}(\mathrm{s})+\mathrm{Na}_{2} \mathrm{S}_{2} \mathrm{O}_{3}(\mathrm{aq}) \rightarrow \mathrm{Na}_{2} \mathrm{S}_{2} \mathrm{O}_{4}(\mathrm{aq})+\mathrm{NaI}(\mathrm{aq})\) (in acid solution)

Determine the oxidation number of the boldface element in each of the following a. \({OF}_{2}\) b. \({UO}_{2}^{2+}\) c. \({RuO}_{4}\) d. \({Fe}_{2} {O}_{3}\)

Identify each of the following changes as either oxidation or reduction. a. \(2 {Cl}^{-} \rightarrow {Cl}_{2}+2 {e}^{-}\) c. \({Ca}^{-2}+2 {e}^{-} \rightarrow 2 {Ca}\) b. \({Na} \rightarrow {Na}^{+}+{e}^{-}\) {d} . \({O}_{2}+4 \mathrm{e}^{-} \rightarrow 2 {O}^{2-}\)

Identify the reducing agents in these equations a. \(4 \mathrm{NH}_{3}+5 \mathrm{O}_{2} \rightarrow 4 \mathrm{NO}+6 \mathrm{H}_{2} \mathrm{O}\) b. \(\mathrm{Na}_{2} \mathrm{SO}_{4}+4 \mathrm{C} \rightarrow \mathrm{Na}_{2} \mathrm{S}+4 \mathrm{CO}\) c. \(4 \mathrm{IrF}_{5}+\mathrm{Ir} \rightarrow 5 \operatorname{Ir} \mathrm{F} 4\)

Write a balanced ionic redox equation using the following pairs of redox half- reactions. a. \(\mathrm{Fe} \rightarrow \mathrm{Fe}^{2+}+2 \mathrm{e}^{-}\) \(\mathrm{Te}^{2+}+2 \mathrm{e}^{-} \rightarrow \mathrm{Te}\) b. IO \(_{4}^{-}+2 \mathrm{e}^{-} \rightarrow \mathrm{IO}_{3}^{-}\) Al \(\rightarrow \mathrm{Al}^{3+}+3 \mathrm{e}^{-}\) (in acid solution) \({c} . {I}_{2}+2 \mathrm{e}^{-} \rightarrow 2 \mathrm{I}^{-}\) \(\mathrm{N}_{2} \mathrm{O} \rightarrow \mathrm{NO}_{3}^{-}+4 \mathrm{e}^{-}(\text { in acid solution })\)

Balance these ionic redox equations by any method. a. \(\mathrm{Sb}^{3+}+\mathrm{MnO}_{4}^{-} \rightarrow \mathrm{SbO}_{4}^{3-}+\mathrm{Mn}^{2+}(\text { in acid solution })\) b. \(\mathrm{N}_{2} \mathrm{O}+\mathrm{ClO}^{-} \rightarrow \mathrm{Cl}^{-}+\mathrm{NO}_{2}^{-}\) (in basic solution)

Gemstones Rubies are gemstones made up mainly of aluminum oxide. Their red color comes from a small amount of chromium(III) ions replacing some of the aluminum ions. Draw the structure of aluminum oxide, and show the reaction in which an aluminum ion is replaced with a chromium ion. Is this a redox reaction?

Balance these ionic redox equations by any method. a. \(\mathrm{Mg}+\mathrm{Fe}^{3+} \rightarrow \mathrm{Mg}^{2+}+\mathrm{Fe}\) b. \(\mathrm{ClO}_{3}^{-}+\mathrm{SO}_{2} \rightarrow \mathrm{Cl}^{-}+\mathrm{SO}_{4}^{2-}(\text { in acid solution })\)

Balance these redox equations by any method. a. \(P+H_{2} O+H N O_{3} \rightarrow H_{3} P O_{4}+N O\) b. \(K C l O_{3}+H C l \rightarrow C l_{2}+C l O_{2}+H_{2} O+K C l\)

Apply The following equations show redox reactions that are sometimes used in the laboratory to generate pure nitrogen gas and pure dinitrogen monoxide gas (nitrous oxide, \(\mathrm{N}_{2} \mathrm{O} )\) $$\mathrm{NH}_{4} \mathrm{NO}_{2}(\mathrm{s}) \rightarrow \mathrm{N}_{2}(\mathrm{g})+2 \mathrm{H}_{2} \mathrm{O}(\mathrm{l})$$ $$\mathrm{NH}_{4} \mathrm{NO}_{3}(\mathrm{s}) \rightarrow \mathrm{N}_{2} \mathrm{O}(\mathrm{g})+2 \mathrm{H}_{2} \mathrm{O}(\mathrm{l})$$ a. Determine the oxidation number of each element in the two equations, and then make diagrams showing the changes in oxidation numbers that occur in each reaction. b. Identify the atom that is oxidized and the atom that is reduced in each of the two reactions. c. Identify the oxidizing and reducing agents in each of the two reactions. d. Write a sentence telling how the electron transfer taking place in these two reactions differs from that taking place here $$2 \mathrm{AgNO}_{3}+\mathrm{Zn} \rightarrow \mathrm{Zn}\left(\mathrm{NO}_{3}\right)_{2}+2 \mathrm{Ag}$$

Predict Consider the fact that all of the following are stable compounds. What can you infer about the oxidation state of phosphorus in its compounds? $$\mathrm{PH}_{3}, \mathrm{PCI}_{3}, \mathrm{P}_{2} \mathrm{H}_{4}, \mathrm{PCI}_{5}, \mathrm{H}_{3} \mathrm{PO}_{4}, \mathrm{Na}_{3} \mathrm{PO}_{3}$$

Solve Potassium permanganate oxidizes chloride ions to chlorine gas. Balance the equation for this redox reaction taking place in acid solution

In the half-reaction \(\mathrm{NO}_{3}^{-} \rightarrow \mathrm{NH}_{4}^{+},\) on which side of the equation should electrons be added? Add the correct number of electrons to the side on which they are needed, and rewrite the equation.

For each reaction described, write the corresponding chemical equation without putting coefficients to balance it. Next, determine the oxidation state of each element in the equation. Then, write the two half-reactions, labeling which is oxidation and which is reduction. Finally, write a balanced equation for the reaction. a. Solid mercuric oxide is put into a test tube and gently heated. Liquid mercury forms on the sides and in the bottom of the tube, and oxygen gas bubbles out from the test tube. b. Solid copper pieces are put into a solution of silver nitrate. Silver metal appears and blue copper(II) nitrate forms in the solution.

A gaseous sample occupies 32.4 mL at ?23°C and 0.75 atm. What volume will it occupy at STP? (Chapter 13)

When iron(III) chloride \(\left(\mathrm{FeCl}_{3}\right)\) reacts in an atmosphere of pure oxygen, the following occurs: $$4 \mathrm{FeCl}_{3}(\mathrm{s})+3 \mathrm{O}_{2}(\mathrm{gv}) \rightarrow 2 \mathrm{Fe}_{2} \mathrm{O}_{3}(\mathrm{s})+6 \mathrm{Cl}_{2}(\mathrm{g})$$ If 45.0 \(\mathrm{g}\) of \(\mathrm{FeCl}_{3}\) reacts and 20.5 \(\mathrm{g}\) of iron(III) oxide is recovered, determine the percent yield. (Chapter 11\()\)

Silverware Practice your technical writing skills by writing a procedure for cleaning tarnished silverware by a redox chemical process. Be sure to include background information describing the process as well as logical steps that would enable anyone to accomplish the task.

Copper was a useful metal even before iron, silver, and gold metals were extracted and used from their ores and used as tools, utensils, jewelry, and artwork. Copper was smelted by heating copper ores with charcoal to high temperatures as early as 8000 years ago. Thousands of pieces of scrap copper have been unearthed in Virginia, where in the 1600s the colonists might have traded this material for food. Compare and contrast the processing and use of copper in those older civilizations with today