World of Chemistry

Steven S.Zumdahl, Susan L.Zumdahl, Donald J.DeCoste

Chapter 18 Oxidation–Reduction Reactions and Electrochemistry - all with Video Answers

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Chapter Questions

01:08

Problem 1

How is oxidation defined? Write an equation showing an atom being oxidized.

Kim Trang Nguyen

Numerade Educator

01:05

Problem 2

How is reduction defined? Write an equation showing an atom being reduced.

Kim Trang Nguyen

Numerade Educator

02:23

For each of the following oxidation–reduction reactions of metals with nonmetals, identify which element is oxidized and which is reduced.
$$
\begin{array}{l}{\text { a. } 6 \mathrm{Na}(s)+\mathrm{N}_{2}(g) \rightarrow 2 \mathrm{Na}_{3} \mathrm{N}(s)} \\ {\text { b. } \mathrm{Mg}(s)+\mathrm{Cl}_{2}(g) \rightarrow \mathrm{MgCl}_{2}(s)} \\ {\text { c. } 2 \mathrm{Al}(s)+3 \mathrm{Br}_{2}(l) \rightarrow 2 \mathrm{AlBr}_{3}(s)} \\ {\text { d. } 4 \mathrm{Fe}(s)+3 \mathrm{O}_{2}(g) \rightarrow 2 \mathrm{Fe}_{2} \mathrm{O}_{3}(s)}\end{array}
$$

Kim Trang Nguyen

Numerade Educator

01:50

Problem 4

For each of the following oxidation–reduction reactions, identify which element is oxidized and which is reduced.
$$
\begin{array}{l}{\text { a. } \mathrm{Mg}(s)+\mathrm{Br}_{2}(l) \rightarrow \mathrm{MgBr}_{2}(s)} \\ {\text { b. } 2 \mathrm{Na}(s)+\mathrm{S}(s) \rightarrow \mathrm{Na}_{2} \mathrm{S}(s)} \\ {\text { c. } \mathrm{Cl}_{2}(g)+2 \mathrm{NaBr}(a q) \rightarrow \mathrm{Br}_{2}(l)+2 \mathrm{NaCl}(a q)} \\ {\text { d. } 6 \mathrm{K}(s)+\mathrm{N}_{2}(g) \rightarrow 2 \mathrm{K}_{3} \mathrm{N}(s)}\end{array}
$$

Kim Trang Nguyen

Numerade Educator

05:01

Problem 5

What is an oxidation state? Why is this concept useful?

Kim Trang Nguyen

Numerade Educator

06:54

Problem 6

Explain why, although it is not an ionic compound, we assign oxygen an oxidation state of $-2$ in water, $\mathrm{H}_{2} \mathrm{O} .$ Give an example of a compound in which oxygen is not in the $-2$ oxidation state.

Kim Trang Nguyen

Numerade Educator

03:46

Problem 7

Why must the sum of all the oxidation states of the atoms in a neutral molecule be zero?

Kim Trang Nguyen

Numerade Educator

01:31

Problem 8

What must be the sum of the oxidation states of all the atoms in a polyatomic ion?

Kim Trang Nguyen

Numerade Educator

03:54

Problem 9

Assign oxidation states to all of the atoms in each of the following:
$$
\begin{array}{ll}{\text { a. } \mathrm{NCl}_{3}} & {\text { c. } \mathrm{PCl}_{5}} \\ {\text { b. } \mathrm{SF}_{6}} & {\text { d. SiH }_{4}}\end{array}
$$

Kim Trang Nguyen

Numerade Educator

03:51

Problem 10

Assign oxidation states to all of the atoms in each of the following:
$$
\begin{array}{ll}{\text { a. HBr }} & {\text { c. } \mathrm{Br}_{2}} \\ {\text { b. HOBr }} & {\text { d. } \mathrm{HBrO}_{4}}\end{array}
$$

Kim Trang Nguyen

Numerade Educator

04:21

Problem 11

Assign oxidation states to all of the atoms in each of the following:
$$
\begin{array}{ll}{\text { a. HNO }_{3}} & {\text { c. HSO }_{4}^{-}} \\ {\text { b. HPO }_{4}^{2-}} & {\text { d. } \mathrm{O}_{2}^{2-}}\end{array}
$$

Kim Trang Nguyen

Numerade Educator

04:36

Problem 12

Assign oxidation states to all of the atoms in each of the following:
$$
\begin{array}{ll}{\text { a. } \mathrm{CuCl}_{2}} & {\text { c. } \mathrm{HCrO}_{4}^{-}} \\ {\text { b. } \mathrm{CrCl}_{3}} & {\text { d. } \mathrm{Cr}_{2} \mathrm{O}_{3}}\end{array}
$$

Kim Trang Nguyen

Numerade Educator

06:26

Problem 13

Assign oxidation states to all of the atoms in each of the following:
$$
\begin{array}{ll}{\text { a. } \mathrm{CH}_{4}} & {\text { c. } \mathrm{KHCO}_{3}} \\ {\text { b. } \mathrm{Na}_{2} \mathrm{CO}_{3}} & {\text { d. } \mathrm{CO}}\end{array}
$$

Kim Trang Nguyen

Numerade Educator

01:42

Problem 14

Oxidation can be defined as a loss of electrons or as an increase in oxidation state. Explain why the two definitions mean the same thing, and give an example to support your explanation.

Nicole Chung

Numerade Educator

01:19

Problem 15

Reduction can be defined as a gain of electrons or as a decrease in oxidation state. Explain why the two definitions mean the same thing, and give an example to support your explanation.

Nicole Chung

Numerade Educator

01:02

Problem 16

Is an oxidizing agent itself oxidized or reduced? Is a reducing agent itself oxidized or reduced?

Ronald Prasad

Numerade Educator

01:42

Problem 17

Does an oxidizing agent increase or decrease its own oxidation state when it acts on another atom? Does a reducing agent increase or decrease its own oxidation state when it acts on another substance?

Abdel Osman

Numerade Educator

03:30

Problem 18

In each of the following reactions, identify which element is oxidized and which is reduced by assigning oxidation numbers.
$$
\begin{array}{l}{\text { a. } \mathrm{Zn}(s)+2 \mathrm{HNO}_{3}(a q) \rightarrow} \\ {\mathrm{Zn}\left(\mathrm{NO}_{3}\right)_{2}(a q)+\mathrm{H}_{2}(g)} \\ {\text { b. } \mathrm{H}_{2}(g)+\mathrm{CuSO}_{4}(a q) \rightarrow \mathrm{Cu}(s)+\mathrm{H}_{2} \mathrm{SO}_{4}(a q)} \\ {\text { c. } \mathrm{N}_{2}(g)+3 \mathrm{Br}_{2}(l) \rightarrow 2 \mathrm{NBr}_{3}(g)} \\ {\text { d. } 2 \mathrm{KBr}(a q)+\mathrm{Cl}_{2}(g) \rightarrow 2 \mathrm{KCl}(a q)+\mathrm{Br}_{2}(l)}\end{array}
$$

Abdel Osman

Numerade Educator

04:27

Problem 19

In each of the following reactions, identify which element is oxidized and which is reduced by assigning oxidation states.
$$
\begin{array}{l}{\text { a. } \mathrm{Cu}(s)+2 \mathrm{AgNO}_{3}(a q) \rightarrow} \\ {2 \mathrm{Ag}(s)+\mathrm{Cu}\left(\mathrm{NO}_{3}\right)_{2}(a q)} \\ {\text { b. } \mathrm{N}_{2}(g)+3 \mathrm{F}_{2}(g) \rightarrow 2 \mathrm{NF}_{3}(g)} \\ {\text { c. } 2 \mathrm{Fe}_{2} \mathrm{O}_{3}(s)+3 \mathrm{S}(s) \rightarrow 4 \mathrm{Fe}(s)+3 \mathrm{SO}_{2}(g)} \\ {\text { d. } 2 \mathrm{H}_{2} \mathrm{O}_{2}(l) \rightarrow 2 \mathrm{H}_{2} \mathrm{O}(l)+\mathrm{O}_{2}(g)}\end{array}
$$

Abdel Osman

Numerade Educator

00:55

Problem 20

In ordinary photography, the light-sensitive portion of the film is a thin coating of a silver halide (usually silver bromide). When the film is exposed to light, the reaction
$$
2 \operatorname{AgBr}(s) \rightarrow 2 \mathrm{Ag}(s)+\mathrm{Br}_{2}(g)
$$
occurs. Identify which element is reduced and which is oxidized.

Abdel Osman

Numerade Educator

01:24

Problem 21

Although magnesium metal does not react with water at room temperature, it does react vigorously with steam at higher temperatures, releasing elemental hydrogen gas from the water.
$$
\mathrm{Mg}(s)+2 \mathrm{H}_{2} \mathrm{O}(g) \rightarrow \mathrm{Mg}(\mathrm{OH})_{2}(s)+\mathrm{H}_{2}(g)
$$
Identify which element is being oxidized and which is being reduced.

Nicole Chung

Numerade Educator

01:59

Problem 22

In what two respects must oxidation–reduction reactions be balanced?

Nailah Jamal

Numerade Educator

01:15

Problem 23

Why must the number of electrons lost in the oxidation equal the number of electrons gained in the reduction? Is it possible to have “leftover” electrons in a reaction?

Nicole Chung

Numerade Educator

03:03

Problem 24

Balance each of the following half-reactions.
$$
\begin{array}{l}{\text { a. } \mathrm{N}_{2}(g) \rightarrow \mathrm{N}_{3}^{-}(a q)} \\ {\text { b. } \mathrm{O}_{2}^{2-}(a q) \rightarrow \mathrm{O}_{2}(g)} \\ {\text { c. } \mathrm{Zn}(s) \rightarrow \mathrm{Zn}^{2+}(a q)} \\ {\text { d. } \mathrm{F}_{2}(g) \rightarrow \mathrm{F}^{-}(a q)}\end{array}
$$

Abdel Osman

Numerade Educator

07:17

Problem 25

Balance each of the following half-reactions, which take place in acidic solution.
$$
\begin{array}{l}{\text { a. } \mathrm{O}_{2}(g) \rightarrow \mathrm{H}_{2} \mathrm{O}(l)} \\ {\text { b. } \mathrm{IO}_{3}^{-}(a q) \rightarrow \mathrm{I}_{2}(s)} \\ {\text { c. } \mathrm{VO}^{2+}(a q) \rightarrow \mathrm{V}^{3+}(a q)} \\ {\text { d. } \mathrm{BiO}^{+}(a q) \rightarrow \mathrm{Bi}(s)}\end{array}
$$

Abdel Osman

Numerade Educator

15:15

Problem 26

Balance each of the following oxidation–reduction reactions, which take place in acidic solution, by using the half-reaction method.
$$
\begin{array}{l}{\text { a. } \mathrm{MnO}_{4}^{-(a q)}+\mathrm{Zn}(s) \rightarrow} \\ {\mathrm{Mn}^{2+}(a q)+\mathrm{Zn}^{2+}(a q)} \\ {\text { b. } \mathrm{Sn}^{4+}(a q)+\mathrm{H}_{2}(g) \rightarrow \mathrm{Sn}^{2+}(a q)+\mathrm{H}^{+}(a q)} \\ {\text { c. } \mathrm{Zn}(s)+\mathrm{NO}_{3}^{-}(a q) \rightarrow \mathrm{Zn}^{2+}(a q)+\mathrm{NO}_{2}(g)} \\ {\text { d. } \mathrm{H}_{2} \mathrm{S}(g)+\mathrm{Br}_{2}(l) \rightarrow \mathrm{S}(s)+\mathrm{Br}^{-}(a q)}\end{array}
$$

Abdel Osman

Numerade Educator

08:05

Problem 27

Iodide ion, $\mathrm{I}^{-},$ is one of the most easily oxidized species. Balance each of the following oxidation-reduction reactions, which take place in acidic solution, by using the half-reaction method.
$$
\begin{array}{l}{\text { a. } \mathrm{IO}_{3}^{-}(a q)+\mathrm{I}^{-}(a q) \rightarrow \mathrm{I}_{2}(a q)} \\ {\text { b. } \mathrm{Cr}_{2} \mathrm{O}_{7}^{2-}(a q)+\mathrm{I}^{-}(a q) \rightarrow} \\ {\mathrm{Cr}^{3+}(a q)+\mathrm{I}_{2}(a q)} \\ {\text { c. } \mathrm{Cu}^{2+}(a q)+\mathrm{I}^{-}(a q) \rightarrow \mathrm{CuI}(s)+\mathrm{I}_{2}(a q)}\end{array}
$$

Leslee Manner

Numerade Educator

02:47

Problem 28

How is an oxidation–reduction reaction set up as a galvanic cell (battery)? How is the transfer of electrons between reducing agent and oxidizing agent made useful?

Gustavo Aroeira

Numerade Educator

03:38

Problem 29

What is a salt bridge? Why is a salt bridge necessary in a galvanic cell? Can a different method be used in place of the salt bridge?

Gustavo Aroeira

Numerade Educator

01:51

Problem 30

In which direction do electrons flow in a galvanic cell, anode to cathode, or vice versa?

Nicole Chung

Numerade Educator

03:28

Problem 31

Consider the oxidation–reduction reaction
$$
\mathrm{Al}(s)+\mathrm{Ni}^{2+}(a q) \rightarrow \mathrm{Al}^{3+}(a q)+\mathrm{Ni}(s)
$$
Sketch a galvanic cell that makes use of this reaction. Which metal ion is reduced? Which metal is oxidized? What half-reaction takes place at the anode in the cell? What half-reaction takes place at the cathode?

Nicole Chung

Numerade Educator

02:55

Problem 32

Consider the oxidation–reduction reaction
$$
\mathrm{Zn}(s)+\mathrm{Pb}^{2+}(a q) \rightarrow \mathrm{Zn}^{2+}(a q)+\mathrm{Pb}(s)
$$
Sketch a galvanic cell that uses this reaction. Which metal ion is reduced? Which metal is oxidized? What half-reaction takes place at the anode in the cell? What half-reaction takes place at the cathode?

Nicole Chung

Numerade Educator

05:17

Problem 33

Write the chemical equation for the overall cell reaction that occurs in a lead storage automobile battery. What species is oxidized in such a battery? What species is reduced? Why can such a battery be “recharged”?

Nicole Chung

Numerade Educator

02:05

Problem 34

Why does an alkaline dry cell battery typically last longer than a normal dry cell? Write the chemical equation for the overall cell reaction in an alkaline dry cell.

Gustavo Aroeira

Numerade Educator

01:14

Problem 35

What process is represented by the corrosion of a metal? Why is corrosion undesirable?

Nicole Chung

Numerade Educator

01:04

Problem 36

Explain how some metals, notably aluminum, naturally resist complete oxidation by the atmosphere.

Nicole Chung

Numerade Educator

01:40

Problem 37

Pure iron ordinarily rusts quickly, but steel does not corrode nearly as fast. How does steel resist corrosion?

Nicole Chung

Numerade Educator

01:33

Problem 38

What is cathodic protection, and how is it applied to prevent oxidation of steel tanks and pipes?

Gustavo Aroeira

Numerade Educator

02:44

Problem 39

What is electrolysis? What types of reactions is electrolysis capable of causing?

Gustavo Aroeira

Numerade Educator

01:22

Problem 40

What reactions go on during the recharging of an automobile battery?

Ronald Prasad

Numerade Educator

01:47

Problem 41

How is aluminum metal produced by electrolysis? Why can’t simpler chemical reduction methods be used?

Gustavo Aroeira

Numerade Educator

01:14

Problem 42

Jewelry is often manufactured by plating an expensive metal such as gold over a cheaper metal. How might such a process be set up as an electrolysis reaction?

Gustavo Aroeira

Numerade Educator

00:48

Problem 43

An oxidizing agent causes the (oxidation/reduction) of another species, and the oxidizing agent itself is (oxidized/reduced).

Ronald Prasad

Numerade Educator

00:35

Problem 44

To function as a good reducing agent, a species must ________ electrons easily.

Ronald Prasad

Numerade Educator

01:22

Problem 45

Jump-starting a dead automobile battery can be dangerous if precautions are not taken, because of the production of an explosive mixture of ___________ and ___________ gases in the battery.

Ronald Prasad

Numerade Educator

03:18

Problem 46

For each of the following unbalanced oxidation–reduction chemical equations, balance the equation by inspection, and identify which species is undergoing oxidation and which is undergoing reduction.
$$
\begin{array}{l}{\text { a. } \mathrm{Fe}(s)+\mathrm{O}_{2}(g) \rightarrow \mathrm{Fe}_{2} \mathrm{O}_{3}(s)} \\ {\text { b. } \mathrm{Al}(s)+\mathrm{Cl}_{2}(g) \rightarrow \mathrm{AlCl}_{3}(s)} \\ {\text { c. } \mathrm{Mg}(s)+\mathrm{P}_{4}(s) \rightarrow \mathrm{Mg}_{3} \mathrm{P}_{2}(s)}\end{array}
$$

Ronald Prasad

Numerade Educator

04:43

Problem 47

Balance each of the following oxidation–reduction reactions, which take place in acidic solution.
$$
\begin{array}{l}{\text { a. } \mathrm{MnO}_{4}^{-}(a q)+\mathrm{H}_{2} \mathrm{O}_{2}(a q) \rightarrow} \\ {\mathrm{Mn}^{2+}(a q)+\mathrm{O}_{2}(g)} \\ {\text { b. } \mathrm{BrO}_{3}^{-}(a q)+\mathrm{Cu}^{+}(a q) \rightarrow} \\ {\mathrm{Br}^{-}(a q)+\mathrm{Cu}^{2+}(a q)} \\ {\text { c. HNO }_{2}(a q)+\mathrm{I}^{-}(a q) \rightarrow \mathrm{NO}(g)+\mathrm{I}_{2}(a q)}\end{array}
$$

Ronald Prasad

Numerade Educator

02:41

Problem 48

For each of the following oxidation–reduction reactions of metals with nonmetals, identify which element is oxidized and which is reduced.
$$
\begin{array}{l}{\text { a. } 3 \mathrm{Zn}(s)+\mathrm{N}_{2}(g) \rightarrow \mathrm{Zn}_{3} \mathrm{N}_{2}(s)} \\ {\text { b. } \mathrm{Co}(s)+\mathrm{S}(s) \rightarrow \cos (s)} \\ {\text { c. } 4 \mathrm{K}(s)+\mathrm{O}_{2}(g) \rightarrow 2 \mathrm{K}_{2} \mathrm{O}(s)} \\ {\text { d. } 4 \mathrm{Ag}(s)+\mathrm{O}_{2}(g) \rightarrow 2 \mathrm{Ag}_{2} \mathrm{O}(s)}\end{array}
$$

Ronald Prasad

Numerade Educator

02:02

Problem 49

Assign oxidation states to all of the atoms in each of the following:
$$
\begin{array}{l}{\text { a. MnO }_{2}} \\ {\text { b. } \mathrm{BaCrO}_{4}} \\ {\text { c. } \mathrm{H}_{2} \mathrm{SO}_{3}} \\ {\text { d. } \mathrm{Ca}_{3}\left(\mathrm{PO}_{4}\right)_{2}}\end{array}
$$

Ronald Prasad

Numerade Educator

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Problem 50

In each of the following reactions, identify which element is oxidized and which is reduced by assigning oxidation states.
$$
\begin{array}{l}{\text { a. } 2 \mathrm{B}_{2} \mathrm{O}_{3}(s)+6 \mathrm{Cl}_{2}(g) \rightarrow 4 \mathrm{BCl}_{3}(l)+3 \mathrm{O}_{2}(g)} \\ {\text { b. GeH }_{4}(g)+\mathrm{O}_{2}(g) \rightarrow \mathrm{Ge}(s)+2 \mathrm{H}_{2} \mathrm{O}(g)} \\ {\text { c. } \mathrm{C}_{2} \mathrm{H}_{4}(g)+\mathrm{Cl}_{2}(g) \rightarrow \mathrm{C}_{2} \mathrm{H}_{4} \mathrm{Cl}_{2}(l)} \\ {\text { d. } \mathrm{O}_{2}(g)+2 \mathrm{F}_{2}(g) \rightarrow 2 \mathrm{OF}_{2}(g)}\end{array}
$$

Ronald Prasad

Numerade Educator

02:21

Problem 51

Consider the oxidation–reduction reaction
$$
\mathrm{Mg}(s)+\mathrm{Cu}^{2+}(a q) \rightarrow \mathrm{Mg}^{2+}(a q)+\mathrm{Cu}(s)
$$
Sketch a galvanic cell that uses this reaction. Which metal ion is reduced? Which metal is oxidized? What half-reaction takes place at the anode in the cell? What half-reaction takes place at the cathode?

Ronald Prasad

Numerade Educator