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Glencoe Chemistry: Matter and Change

Buthelezi ,Dingrando,Wistrom,Zike

Chapter 19

Redox Reactions

Educators


Problem 1

Identify each of the following changes as either oxidation or reduction.
Recall that $\mathrm{e}^{-}$ is the symbol for an electron.
$$
\begin{array}{ll}{\text { a. } 1_{2}+2 \mathrm{e}^{-} \rightarrow 21^{-}} & {\text { c. } \mathrm{Fe}^{2+} \rightarrow \mathrm{Fe}^{3+}+\mathrm{e}^{-}} \\ {\text { b. } \mathrm{K} \rightarrow \mathrm{K}^{+}+\mathrm{e}^{-}} & {\text { d. } \mathrm{Ag}^{+}+\mathrm{e}^{-} \rightarrow \mathrm{Ag}}\end{array}
$$

Stephen P.
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Problem 2

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

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Numerade Educator

Problem 3

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

Stephen P.
Numerade Educator

Problem 4

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

Melissa G.
Numerade Educator

Problem 5

Determine the oxidation number of the boldface element in the following formulas for
compounds.
$$
\text { a. } \mathrm{NaClO}_{4} \quad \text { b. } \mathrm{AlPO}_{4} \quad \text {c} . \mathrm{HNO}_{2}
$$

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Problem 6

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

Melissa G.
Numerade Educator

Problem 7

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

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Problem 8

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

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Problem 9

Explain why oxidation and reduction must always occur together.

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Problem 10

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

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Problem 11

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

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Problem 12

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

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Numerade Educator

Problem 13

Determine the oxidation number of the boldface element in these ions.
$$
\begin{array}{ll}{\text { a. } 10_{4}^{-}} & {\text { c. } B_{4} 0_{7}^{2-}} \\ {\text { b. } M n O_{4}-} & {\text { d. } N H_{2}-}\end{array}
$$

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Problem 14

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.

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Problem 15

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

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Problem 16

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

Melissa G.
Numerade Educator

Problem 17

Use the oxidation-number method to balance these redox equations.
$$
\mathrm{NH}_{3}(\mathrm{g})+\mathrm{NO}_{2}(\mathrm{g}) \rightarrow \mathrm{N}_{2}(\mathrm{g})+\mathrm{H}_{2} \mathrm{O}(1)
$$

Stephen P.
Numerade Educator

Problem 18

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

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Problem 19

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

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Problem 20

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

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Problem 21

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

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Problem 22

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

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Problem 23

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.
$$
\mathrm{Cr}_{2} \mathrm{O}_{7}^{-}(\mathrm{aq})+\mathrm{I}^{-}(\mathrm{aq}) \rightarrow \mathrm{Cr}^{3+}(\mathrm{aq})+\mathrm{I}_{2}(\mathrm{s}) \quad \text {(in acid solution)}
$$

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Numerade Educator

Problem 24

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.
$$
\mathrm{Mn}^{2+}(\mathrm{aq})+\mathrm{BiO}_{3}^{-}(\mathrm{aq}) \rightarrow \mathrm{MnO}_{4}^{-}(\mathrm{aq})+\mathrm{Bi}^{2+}(\mathrm{aq}) \quad \text {(in acid solution)}
$$

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Problem 25

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.
$$
\text {Challenge} \quad \mathrm{N}_{2} \mathrm{O}(\mathrm{g})+\mathrm{ClO}^{-}(\mathrm{aq}) \rightarrow \mathrm{NO}_{2}^{-}(\mathrm{aq})+\mathrm{Cl}^{-}(\mathrm{aq}) \quad \text {(in basic solution)}
$$

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Problem 26

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?

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Problem 27

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.

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Problem 28

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

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Problem 29

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

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Problem 30

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

Melissa G.
Numerade Educator

Problem 31

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?

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Problem 32

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

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Problem 33

What is the main characteristic of oxidation-reduction
reactions?

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Problem 34

Explain why not all oxidation reactions involve oxygen.

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Problem 35

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

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Problem 36

Define oxidation number.

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Problem 37

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

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Problem 38

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?

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Problem 39

What is the oxidation number for copper in each of the
compounds shown in Figure 19.9$?$

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Problem 40

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.

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Problem 41

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

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Problem 42

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

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Numerade Educator

Problem 43

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

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Problem 44

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

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Problem 45

Determine the oxidation number of the boldface element in these substances and ions.
$$
\begin{array}{ll}{\text { a. } \mathrm{CaCrO}_{4}} & {\text { c. } N O_{2}-} \\ {\text { b. } N a H S O_{4}} & {\text { d. } B r O_{3}^{-}}\end{array}
$$

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Problem 46

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

Melissa G.
Numerade Educator

Problem 47

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

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Numerade Educator

Problem 48

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

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Problem 49

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

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Problem 50

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

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Problem 51

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

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Problem 52

Explain why writing hydrogen ions as $\mathrm{H}^{+}$ in redox
reactions represents a simplification and not how
they exist.

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Problem 53

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?

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Problem 54

Explain what a spectator ion is.

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Problem 55

Define the term species in terms of redox reactions.

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Problem 56

Is the following equation balanced? Explain.
$$
\mathrm{Fe}(\mathrm{s})+\mathrm{Ag}^{+}(\mathrm{aq}) \rightarrow \mathrm{Fe}^{2+}(\mathrm{aq})+\mathrm{Ag}(\mathrm{s})
$$

Melissa G.
Numerade Educator

Problem 57

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

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Problem 58

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

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Problem 59

Use the oxidation-number method to balance these redox equations.
\begin{equation}
\begin{array}{l}{\text { a. } \mathrm{Cl}_{2}+\mathrm{NaOH} \rightarrow \mathrm{NaCl}+\mathrm{HOCl}} \\ {\text { b. } \mathrm{HBrO}_{3} \rightarrow \mathrm{Br}_{2}+\mathrm{H}_{2} \mathrm{O}+\mathrm{O}_{2}}\end{array}
\end{equation}

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Problem 60

Balance these net ionic equations for redox reactions.
\begin{equation}
\begin{array}{l}{\text { a. } \mathrm{Au}^{3+}(\mathrm{aq})+\mathrm{I}^{-}(\mathrm{aq}) \rightarrow \mathrm{Au}(\mathrm{s})+\mathrm{I}_{2}(\mathrm{s})} \\ {\text { b. } \mathrm{Ce}^{4+}(\mathrm{aq})+\mathrm{Sn}^{2+}(\mathrm{aq}) \rightarrow \mathrm{Ce}^{3+}(\mathrm{aq})+\mathrm{Sn}^{4+}(\mathrm{aq})}\end{array}
\end{equation}

Melissa G.
Numerade Educator

Problem 61

Use the oxidation-number method to balance the following ionic redox equations.
\begin{equation}
\begin{array}{l}{\text { a. } \mathrm{Al}+\mathrm{I}_{2} \rightarrow \mathrm{Al}^{3+}+\mathrm{I}^{-}} \\ {\text { b. } \mathrm{MnO}_{2}+\mathrm{Br}^{-} \rightarrow \mathrm{Mn}^{2+}+\mathrm{Br}_{2}(\text { in acid solution })}\end{array}
\end{equation}

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Problem 62

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

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Problem 63

Sapphire The mineral corundum is comprised of aluminum oxide $\left(\mathrm{Al}_{2} \mathrm{O}_{3}\right)$ and is colorless. Sapphire is mostly aluminum oxide, but it contains small amounts of $\mathrm{Fe}^{2+}$ and $\mathrm{Ti}^{4+} .$ The color of sapphire results from an electron transfer from $\mathrm{Fe}^{2+}$ to $\mathrm{Ti}^{4+} .$ Based on Figure $19.11,$ draw
the reaction that occurs resulting in the mineral on the right. What are the oxidizing and reducing agents?

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Problem 64

Write the oxidation and reduction half-reactions represented in each of these redox equations. Write the half-reactions in net ionic form if they occur in aqueous solution.
\begin{equation}
\begin{array}{l}{\text { a. } \mathrm{PbO}(\mathrm{s})+\mathrm{NH}_{3}(\mathrm{g}) \rightarrow \mathrm{N}_{2}(\mathrm{g})+\mathrm{H}_{2} \mathrm{O}(1)+\mathrm{Pb}(\mathrm{s})} \\ {\text { 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})} \\ {\text { c. } \mathrm{Sn}(\mathrm{s})+2 \mathrm{HCl}(\mathrm{aq}) \rightarrow \mathrm{SnCl}_{2}(\mathrm{aq})+\mathrm{H}_{2}(\mathrm{g})}\end{array}
\end{equation}

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Problem 65

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

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Problem 66

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

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Problem 67

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.

Stephen P.
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Problem 68

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

Melissa G.
Numerade Educator

Problem 69

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.
\begin{equation}
\begin{array}{l}{\text { a. } \mathrm{NH}_{3}(\mathrm{g})+\mathrm{NO}_{2}(\mathrm{g}) \rightarrow \mathrm{N}_{2}(\mathrm{g})+\mathrm{H}_{2} \mathrm{O}(1)} \\ {\text { b. } \mathrm{Br}_{2} \rightarrow \mathrm{Br}^{-}+\mathrm{BrO}_{3}^{-} \text { (in basic solution) }}\end{array}
\end{equation}

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Problem 70

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.
$$\begin{array}{l}{\mathrm{KMnO}_{4}(\mathrm{aq})+\mathrm{FeSO}_{4}(\mathrm{aq})+\mathrm{H}_{2} \mathrm{SO}_{4}(\mathrm{aq}) \rightarrow} \\ {\quad \mathrm{Fe}_{2}\left(\mathrm{SO}_{4}\right)_{3}(\mathrm{aq})+\mathrm{MnSO}_{4}(\mathrm{aq})+} \\ \quad {\mathrm{K}_{2} \mathrm{SO}_{4}(\mathrm{aq})+\mathrm{H}_{2} \mathrm{O}(\mathrm{l})}\end{array}$$

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Problem 71

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.
\begin{equation}
\begin{array}{l}{\text { a. } \mathrm{PbO}(\mathrm{s})+\mathrm{NH}_{3}(\mathrm{g}) \rightarrow \mathrm{N}_{2}(\mathrm{g})+\mathrm{H}_{2} \mathrm{O}(1)+\mathrm{Pb}(\mathrm{s})} \\ {\text { b. } \mathrm{I}_{2}(\mathrm{s})+\mathrm{NaS}_{2} \mathrm{O}_{3}(\mathrm{aq}) \rightarrow \mathrm{Na}_{2} \mathrm{S}_{2} \mathrm{O}_{4}(\mathrm{aq})+\mathrm{NaI}(\mathrm{aq})} \\ {\text { c. } \mathrm{Sn}(\mathrm{s})+2 \mathrm{HCl}(\mathrm{aq}) \rightarrow \mathrm{SnCl}_{2}(\mathrm{aq})+\mathrm{H}_{2}(\mathrm{g})}\end{array}
\end{equation}

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Problem 72

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.
\begin{equation}
\begin{array}{l}{\text { a. } \mathrm{Cl}^{-}(\mathrm{aq})+\mathrm{NO}_{3}^{-}(\mathrm{aq}) \rightarrow \mathrm{ClO}^{-}(\mathrm{aq})+\mathrm{NO}(\mathrm{g})} \\ {\text { (in acid solution) }} \\ {\text { b. } \mathrm{IO}_{3}-(\mathrm{aq})+\mathrm{Br}^{-}(\mathrm{aq}) \rightarrow \mathrm{Br}_{2}(1)+\mathrm{IBr}(\mathrm{s}) \\ {\text { (in acid solution) }}} \\ {\text { 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})} \\ {(\text { in acid solution })} \end{array}
\end{equation}

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Problem 73

Determine the oxidation number of the boldface element in each of the following.
\begin{equation}
\quad \text { a. }\mathrm{OF}_{2} \quad \text { b. UO }_{2}^{2+} \quad \text { c. } \operatorname{Ru} \mathrm{O}_{4} \quad \text { d. Fe }_{2} \mathrm{O}_{3}
\end{equation}

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Problem 74

Identify each of the following changes as either oxidation or reduction.
\begin{equation}
\begin{array}{ll}{\text { a. } 2 \mathrm{Cl}^{-} \rightarrow \mathrm{Cl}_{2}+2 \mathrm{e}^{-}} & {\text { c. } \mathrm{Ca}^{-2}+2 \mathrm{e}^{-} \rightarrow 2 \mathrm{Ca}} \\ {\text { b. } \mathrm{Na} \rightarrow \mathrm{Na}^{+}+\mathrm{e}^{-}} & {\text { d. } \mathrm{O}_{2}+4 \mathrm{e}^{-} \rightarrow 2 \mathrm{O}^{2-}}\end{array}
\end{equation}

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Problem 75

Use the rules for assigning oxidation numbers to complete Table $19.7 .$

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Problem 76

Identify the reducing agents in these equations.
\begin{equation}
\begin{array}{l}{\text { a. } 4 \mathrm{NH}_{3}+5 \mathrm{O}_{2} \rightarrow 4 \mathrm{NO}+6 \mathrm{H}_{2} \mathrm{O}} \\ {\text { b. } \mathrm{Na}_{2} \mathrm{SO}_{4}+4 \mathrm{C} \rightarrow \mathrm{Na}_{2} \mathrm{S}+4 \mathrm{CO}} \\ {\text { c. } 4 \mathrm{IrF}_{5}+\mathrm{Ir} \rightarrow 5 \mathrm{IrF} 4}\end{array}
\end{equation}

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Problem 77

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

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Problem 78

What is the oxidation number of chromium in each of the compounds shown in Figure 19.13$?$

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Problem 79

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

Stephen P.
Numerade Educator

Problem 80

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?

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Problem 81

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

Stephen P.
Numerade Educator

Problem 82

Balance these redox equations by any method.
\begin{equation}
\begin{array}{l}{\text { a. } \mathrm{P}+\mathrm{H}_{2} \mathrm{O}+\mathrm{HNO}_{3} \rightarrow \mathrm{H}_{3} \mathrm{PO}_{4}+\mathrm{NO}} \\ {\text { b. } \mathrm{KClO}_{3}+\mathrm{HCl} \rightarrow \mathrm{Cl}_{2}+\mathrm{ClO}_{2}+\mathrm{H}_{2} \mathrm{O}+\mathrm{KCl}}\end{array}
\end{equation}

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Problem 83

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, $\mathrm{N}_{2} \mathrm{O} )$.
$$\begin{array}{c}{\mathrm{NH}_{4} \mathrm{NO}_{2}(\mathrm{s}) \rightarrow \mathrm{N}_{2}(\mathrm{g})+2 \mathrm{H}_{2} \mathrm{O}(1)} \\ {\mathrm{NH}_{4} \mathrm{NO}_{3}(\mathrm{s}) \rightarrow \mathrm{N}_{2} \mathrm{O}(\mathrm{g})+2 \mathrm{H}_{2} \mathrm{O}(\mathrm{l})}\end{array}$$
\begin{equation}
\begin{array}{l}{\text { a. Determine the oxidation number of each element in }} \\ {\text { the two equations, and then make diagrams showing }} \\ {\text { the changes in oxidation numbers that occur in each }} \\ {\text { reaction. }} \\ {\text { b. Identify the atom that is oxidized and the atom that is }} \\ {\text { reduced in each of the two reactions. }} \\ {\text { c. Identify the oxidizing and reducing agents in each of }} \\ {\text { the two reactions. }} \\ {\text { d. Write a sentence telling how the electron transfer }} \\ {\text { taking place in these two reactions differs from }} \\ {\text { that taking place here. }} \end{array}\end{equation}
$$2 \mathrm{AgNO}_{3}+\mathrm{Zn} \rightarrow \mathrm{Zn}\left(\mathrm{NO}_{3}\right)_{2}+2 \mathrm{Ag}$$

Stephen P.
Numerade Educator

Problem 84

Analyze Examine the net ionic equation below for the reaction that occurs when the thiosulfate ion $\left(\mathrm{S}_{2} \mathrm{O}_{3}^{2-}\right)$ is oxidized to the tetrathionate ion $\left(\mathrm{S}_{4} \mathrm{O}_{6}^{2-}\right) .$ Balance the equation using the half-reaction method. Figure 19.14 will help you to determine the oxidation numbers to use.
$$\mathrm{S}_{2} \mathrm{O}_{3}^{2-}+\mathrm{I}_{2} \rightarrow \mathrm{I}^{-}+\mathrm{S}_{4} \mathrm{O}_{6}^{2-}(\text { in acid solution })$$

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Problem 85

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?
$$\mathrm{PH}_{3}, \mathrm{PCl}_{3}, \mathrm{P}_{2} \mathrm{H}_{4}, \mathrm{PCI}_{5}, \mathrm{H}_{3} \mathrm{PO}_{4}, \mathrm{Na}_{3} \mathrm{PO}_{3}$$

Stephen P.
Numerade Educator

Problem 86

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

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Problem 87

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.

Stephen P.
Numerade Educator

Problem 88

The redox reaction between dichromate ion and iodide ion in acid solution is shown in Figure $19.15 .$ Use the half-reaction method to balance the equation for this redox reaction.

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Problem 89

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.
\begin{equation}
\begin{array}{l}{\text { a. Solid mercuric oxide is put into a test tube and gently }} \\ {\text { heated. Liquid mercury forms on the sides and in the }} \\ {\text { bottom of the tube, and oxygen gas bubbles out from }} \\ {\text { the test tube. }} \\ {\text { b. Solid copper pieces are put into a solution of silver }} \\ {\text { nitrate. Silver metal appears and blution of silver }} \\ {\text { nitrate forms in the solution. }}\end{array}
\end{equation}

Stephen P.
Numerade Educator

Problem 90

A gaseous sample occupies 32.4 $\mathrm{mL}$ at $-23^{\circ} \mathrm{C}$ and 0.75 atm. What volume will it occupy at $\mathrm{STP}$? (Chapter 13$)$

Melissa G.
Numerade Educator

Problem 91

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 g of $\mathrm{FeCl}_{3}$ reacts and 20.5 g of iron(III) oxide is recovered, determine the percent yield. (Chapter 11$)$

Stephen P.
Numerade Educator

Problem 92

Steel Research the role of oxidation-reduction reactions in the manufacture of steel. Write a summary of your findings, including appropriate diagrams and equations representing the reactions.

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Problem 93

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.

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Problem 94

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 1600 s the colonists might have traded this material for food. Compare and contrast the processing and use of copper in those older civilizations with today.

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Problem 95

Write the equation for what has occurred in the pottery shown in Figure $19.16 .$

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Problem 96

Based on the color of the pottery, what is the oxidation state of the copper that is reduced? Oxidized?

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