Chapter 19

Identify each of the following changes as either oxidation or reduction. Recall that \(e^{-}\) is the symbol for an electron. \(\begin{array}{ll}{\text { a. I} _{2}+2 \mathrm{e}^{-} \rightarrow 2\text{I}^{-}} & {\text { c. } \mathrm{Fe}^{2+} \rightarrow \mathrm{Fe}^{3+}+\mathrm{e}^{-}} \\ {\text { b. } \mathrm{K} \rightarrow \mathrm{K}^{+}+\mathrm{e}^{-}} & {\text { d. } \mathrm{Ag}^{+}+\mathrm{e}^{-} \rightarrow \mathrm{Ag}}\end{array}\)

Identify what is oxidized and what is reduced in the following processes. $$ \begin{array}{l}{\text { a. } 2 \mathrm{Br}^{-}+\mathrm{Cl}_{2} \rightarrow \mathrm{Br}_{2}+2 \mathrm{Cl}^{-}} \\ {\text { b. } 2 \mathrm{Ce}+3 \mathrm{Cu}^{2+} \rightarrow 3 \mathrm{Cu}+2 \mathrm{Ce}^{3+}} \\ {\text { c. } 2 \mathrm{zn}+\mathrm{O}_{2} \rightarrow 2 \mathrm{nO}} \\ {\text { d. } 2 \mathrm{Na}+2 \mathrm{H}^{+} \rightarrow 2 \mathrm{Na}^{+}+\mathrm{H}_{2}}\end{array} $$

Identify the oxidizing agent and the reducing agent in the following equation. Explain your answer. $$ \mathrm{Fe}(\mathrm{s})+\mathrm{Ag}+(\mathrm{aq}) \rightarrow \mathrm{Fe}^{2+}(\mathrm{aq})+\mathrm{Ag}(\mathrm{s}) $$

Challenge Identify the oxidizing agent and the reducing agent in each reaction. $$ \begin{array}{l}{\text { a. } M g+I_{2} \rightarrow M g l_{2}} \\ {\text { b. } H_{2} S+C l_{2} \rightarrow S+2 H C l}\end{array} $$

Determine the oxidation number of the boldface element in the following formulas for compounds. a. Na\(Cl\)O\(_{4}\) b. Al\(P\)O\(_4\) c. H\(N\)O\(_{2}\)

Determine the oxidation number of the boldface element in the following formulas for ions. $$ \text { a. } \mathrm{NH}_{4}^+ \quad \text { b. } \mathrm{AsO}_{4}^{3-} \quad \text { c. } \mathrm{CrO}_{4}^{2-} $$

Determine the oxidation number of nitrogen in each of these molecules or ions. $$ \text { a. } \mathrm{NH}_{3} \quad \text { b. KCN } \quad \text { c. } \mathrm{N}_{2} \mathrm{H}_{4} $$

Challenge Determine the net change of oxidation number of each of the elements in these redox equations. $$ \begin{array}{l}{\text { a. } \mathrm{C}+\mathrm{O}_{2} \rightarrow \mathrm{CO}_{2}} \\ {\text { b. } \mathrm{Cl}_{2}+\mathrm{Znl}_{2} \rightarrow \mathrm{Znl}_{2}+\mathrm{I}_{2}} \\ {\text { c. } \mathrm{CdO}+\mathrm{CO} \rightarrow \mathrm{Cd}+\mathrm{CO}_{2}}\end{array} $$

Explain why oxidation and reduction must always occur together.

Describe the roles of oxidizing agents and reducing agents in a redox reaction. How is each changed in the reaction?

The equation for the reaction of iron metal with hydrobromic acid to form iron(III) bromide and hydrogen gas. Determine the net change in oxidation for the element that is reduced and the element that is oxidized.

Determine the oxidation number of the boldface element in these compounds. $$ \begin{array}{ll}{\text { a. HNO }_{3}} & {\text { c. Sb}_{2} \text {O}_{5}} \\\ {\text { b. CaN}_{2}} & {\text { d. CuWO }_{4}}\end{array} $$

Determine the oxidation number of the boldface element in these ions. $$ \begin{array}{ll}{\text { a. IO} _{4}^{-}} & {\text { c. } B_{4} 0_{7}^{2-}} \\\ {\text { b. MnO} _{4}-} & {\text { d. NH} _{2}-}\end{array} $$

Make and Use Graphs Alkali metals are strong reducing agents. Make a graph showing how the reducing abilities of the alkali metals wouls increase or decrease as you move down the family from sodium to francium.

Use the oxidation-number method to balance this redox equations. $$ \mathrm{HCl}+\mathrm{HNO}_{3} \rightarrow \mathrm{HOCl}+\mathrm{NO}+\mathrm{H}_{2} \mathrm{O} $$

Use the oxidation-number method to balance these redox equations. $$ \mathrm{SnCl}_{4}+\mathrm{Fe} \rightarrow \mathrm{SnCl}_{2}+\mathrm{FeCl}_{3} $$

Use the oxidation-number method to balance these redox equations. $$ \mathrm{NH}_{3}(\mathrm{~g})+\mathrm{NO}_{2}(\mathrm{~g}) \rightarrow \mathrm{N}_{2}(\mathrm{~g})+\mathrm{H}_{2} \mathrm{O}(\mathrm{I}) $$

Use the oxidation-number method to balance these redox equations. $$ \text {Challenge} \quad \mathrm{SO}_{2}+\mathrm{Br}_{2}+\mathrm{H}_{2} \mathrm{O} \rightarrow \mathrm{HBr}+\mathrm{H}_{2} \mathrm{SO}_{4} $$

Use the oxidation-number method to balance the following net ionic redox equations. $$ \mathrm{H}_{2} \mathrm{S}(\mathrm{g})+\mathrm{NO}_{3}^{-}(\mathrm{aq}) \rightarrow \mathrm{S}(\mathrm{s})+\mathrm{NO}(\mathrm{g}) \quad \text {(in acid solution)} $$

Use the oxidation-number method to balance the following net ionic redox equations. $$ \mathrm{Cr}_{2} \mathrm{O}_{7}^{2-}(\mathrm{aq})+\mathrm{I}^{-}(\mathrm{aq}) \rightarrow \mathrm{Cr}^{3+}(\mathrm{aq})+\mathrm{I}_{2}(\mathrm{s}) \quad \text {(in acid solution)} $$

Use the oxidation-number method to balance the following net ionic redox equations. $$ \mathrm{Zn}+\mathrm{NO}_{3}^{-} \rightarrow \mathrm{Zn}^{2+}+\mathrm{NO}_{2} \quad \text {(in acid solution)} $$

Use the oxidation-number method to balance the following net ionic redox equations. $$ \text {Challenge} \quad \mathrm{I}-(\mathrm{aq})+\mathrm{MnO}_{4}^{-}(\mathrm{aq}) \rightarrow \mathrm{I}_{2}(\mathrm{s})+\mathrm{MnO}_{2}(\mathrm{s}) \quad \text {(in basic solution)} $$

Use the half-reaction method to balance the redox equations. Begin by writing the oxidation and reduction half-reactions. Leave the balanced equation in ionic form. $$ \mathrm{Cr}_{2} \mathrm{O}_{7}^{-}(\mathrm{aq})+\mathrm{I}^{-}(\mathrm{aq}) \rightarrow \mathrm{Cr}^{3+}(\mathrm{aq})+\mathrm{I}_{2}(\mathrm{s}) \quad \text {(in acid solution)} $$

Use the half-reaction method to balance the redox equations. Begin by writing the oxidation and reduction half-reactions. Leave the balanced equation in ionic form. $$ \mathrm{Mn}^{2+}(\mathrm{aq})+\mathrm{BiO}_{3}^{-}(\mathrm{aq}) \rightarrow \mathrm{MnO}_{4}^{-}(\mathrm{aq})+\mathrm{Bi}^{2+}(\mathrm{aq}) \quad \text {(in acid solution)} $$

Use the half-reaction method to balance the redox equations. Begin by writing the oxidation and reduction half-reactions. Leave the balanced equation in ionic form. $$ \text {Challenge} \quad \mathrm{N}_{2} \mathrm{O}(\mathrm{g})+\mathrm{ClO}^{-}(\mathrm{aq}) \rightarrow \mathrm{NO}_{2}^{-}(\mathrm{aq})+\mathrm{Cl}^{-}(\mathrm{aq}) \quad \text {(in basic solution)} $$

Explain how changes in oxidation number are related to the electrons transferred in a redox reaction. How are the changes related to the processes of oxidation and reduction?

Describe why it is important to know the conditions under which an aqueous oxidation-reducation reaction takes place in order to balance the ionic equation for the reaction.

Explain the steps of the oxidation-number method of balancing equations.

State what an oxidation half-reaction shows. What does a reduction half- reaction show?

Write the oxidation and reduction half-reactions for the redox equation. $$ \mathrm{Pb}(\mathrm{s})+\mathrm{Pd}\left(\mathrm{NO}_{3}\right)_{2}(\mathrm{aq}) \rightarrow \mathrm{Pb}\left(\mathrm{NO}_{3}\right)_{2}(\mathrm{aq})+\mathrm{Pd}(\mathrm{s}) $$

Determine The oxidation half-reaction of a redox reaction is \(\mathrm{Sn}^{2+} \rightarrow \mathrm{Sn}^{4+}+2 \mathrm{e}^{-},\) and the reduction half-reaction is \(\mathrm{Au}^{3+}+3 \mathrm{e}^{-} \rightarrow\) Au. What minimum numbers of tin(ll) ions and gold(ll) ions would have to react in order to have zero electrons left over?

Apply Balance the following equations. $$ \begin{array}{l}{\text { a. } \mathrm{HClO}_{3}(\mathrm{aq}) \rightarrow \mathrm{ClO}_{2}(\mathrm{g})+\mathrm{HClO}_{4}(\mathrm{aq})+\mathrm{H}_{2} \mathrm{O}(\mathrm{l})} \\ {\text { b. } \mathrm{H}_{2} \mathrm{SeO}_{3}(\mathrm{aq})+\mathrm{HClO}_{3}(\mathrm{aq}) \rightarrow \mathrm{H}_{2} \mathrm{SeO}_{4}(\mathrm{aq})+\mathrm{Cl}_{2}(\mathrm{g})+\mathrm{H}_{2} \mathrm{O}(\mathrm{l})} \\ \text { c. } \mathrm{Cr}_{2} \mathrm{O}_{7}^{2-}(\mathrm{aq})+\mathrm{Fe}^{2+}(\mathrm{aq}) \rightarrow \mathrm{Cr}^{3+}(\mathrm{aq})+\mathrm{Fe}^{3+}(\mathrm{aq}) \quad \text {(in acid solution)}\end{array} $$

What is the main characteristic of oxidation-reduction reactions?

Explain why not all oxidation reactions involve oxygen.

In terms of electrons, what happens when an atom is oxidized? When an atom is reduced?

Define oxidation number.

Metals What is the oxidation number of alkaline earth metals in their compounds? Of alkali metals?

How does the oxidation number in an oxidation process relate to the number of electrons lost? How does the change in oxidation number in a reduction process relate to the number of electrons gained?

Copper and air Copper statues, such as the Statue of Liberty, begin to appear green after they have been exposed to air. In this redox process, copper metal reacts with oxygen to form solid copper oxide, which forms the green coating. Write the reaction for this redox process, and identify what is oxidized and what is reduced in the process.

Identify the species oxidized and the species reduced in each of these redox equations. $$ \begin{array}{l}{\text { a. } 3 \mathrm{Br}_{2}+2 \mathrm{Ga} \rightarrow 2 \mathrm{GaBr}_{3}} \\ {\text { b. } \mathrm{HCl}+\mathrm{Zn} \rightarrow \mathrm{ZnCl}_{2}+\mathrm{H}_{2}} \\ {\text { c. } \mathrm{Mg}+\mathrm{N}_{2} \rightarrow \mathrm{Mg}_{3} \mathrm{N}_{2}}\end{array} $$

Identify the oxidizing agent and the reducing agent in each of these redox equations. $$ \begin{array}{l}{\text { a. } \mathrm{N}_{2}+3 \mathrm{H}_{2} \rightarrow 2 \mathrm{NH}_{3}} \\ {\text { b. } 2 \mathrm{Na}+\mathrm{I}_{2} \rightarrow 2 \mathrm{NaI}}\end{array} $$

What is the reducing agent in this balanced equation? $$ \begin{array}{c}{8 \mathrm{H}^{+}+\mathrm{Sn}+6 \mathrm{Cl}^{-}+4 \mathrm{NO}_{3}^{-1} \rightarrow} \\ {\mathrm{SnCl}_{6}^{-2}+4 \mathrm{NO}_{2}+4 \mathrm{H}_{2} \mathrm{O}}\end{array} $$

What is the oxidation number of manganese in \(\mathrm{KMnO}_{4} ?\)

Determine the oxidation number of the boldface element in these substances and ions. a. \(\mathrm{CaCrO}_{4}\) b. \(\mathrm{NaHSO}_{4}\) c. \(\mathrm{NO}_{2}^{-}\) d. \(\mathrm{BrO}_{3}^{-}\)

Identify each of these half-reactions as either oxidation or reduction. $$ \begin{array}{l}{\text { a. } \mathrm{Al} \rightarrow \mathrm{Al}^{3+}+3 \mathrm{e}^{-}} \\ {\text { b. } \mathrm{Cu}^{2+}+\mathrm{e}^{-} \rightarrow \mathrm{Cu}^{+}}\end{array} $$

Which of these equations does not represent a redox reaction? Explain your answer. $$ \begin{array}{l}{\text { a. } \mathrm{LiOH}+\mathrm{HNO}_{3} \rightarrow \mathrm{LiNO}_{3}+\mathrm{H}_{2} \mathrm{O}} \\ {\text { b. } \mathrm{MgI}_{2}+\mathrm{Br}_{2} \rightarrow \mathrm{MgBr}_{2}+\mathrm{I}_{2}}\end{array} $$

Determine the oxidation number of nitrogen in each of these molecules or ions. $$ \text { a. }\mathrm{NO}_{3} \quad \text { b. } \mathrm{N}_{2} \mathrm{O} \quad \text { c. } \mathrm{NF}_{3} $$

Determine the oxidation number of each element in these compounds or ions. $$ \begin{array}{l}{\text { a. } \mathrm{Au}_{2}\left(\mathrm{SeO}_{4}\right)_{3} \text { (gold (III) selenate) }} \\ {\text { b. } \mathrm{Ni}(\mathrm{CN})_{2} \text { (nickel (II) cyanide) }}\end{array} $$

Explain how the sulfite ion \(\left(5 \mathrm{O}_{3}^{2-}\right)\) differs from sulfur trioxide \(\left(\mathrm{SO}_{3}\right),\) shown in Figure \(19.10 .\)

Compare and contrast balancing redox equations in acidic and basic solutions.