Kelter P., Mosher M., Scott A. Chemistry. The Practical Science

Подождите немного. Документ загружается.

is the isolation of pure metals from a solution of metal ions, and it has been

known for quite some time. Humphrey Davey, an English chemist, used this

process to isolate metallic sodium from NaCl in 1807. However, the electrowin-

ning of aluminum wasn’t invented until 1886. Working independently, Charles

M. Hall (an American chemist) and Paul Heroult (a French chemist) discovered

that if bauxite ore was puriﬁed to alumina (Al

) and dissolved in molten cryo-

lite (Na

AlF

), metallic aluminum could be made. In the Hall–Heroult process,an

electrical current is passed into a molten mixture of alumina and cryolite to make

molten aluminum (Figure 19.19). The overall reaction used in the Hall–Heroult

process, shown below, is much simpler than the reactions that take place at the

electrodes. Why can’t an electric current be applied directly to an aqueous solu-

tion of Al

in order to produce aluminum metal? Because the potential is so

large, instead of obtaining aluminum metal from the aqueous solution, the water

undergoes electrolysis to form hydrogen and oxygen gases.

2Al

(l) + 3C(s) n 4Al(l) + 3CO

(g) E° ≈−2.1 V

The negative potential tells us that this redox reaction is not spontaneous.

In fact, the reverse reaction—that is, 4Al(l) + 3CO

(g) n 2Al

(l) + 3C(s)—

is quite spontaneous (E

° ≈

+2.1 V). Modifying the concentrations and tempera-

tures helps a little to make the potential of the overall reaction less negative, via

the Nernst equation, but not enough to make the reaction spontaneous. If we

apply a positive potential to the reaction that is larger than the negative potential

expressed by the electrochemical cell, we can force the reaction to go forward.

Examination of a simpler redox reaction can be helpful here. For example, in

concept, if we wish to make the copper–iron redox reaction spontaneous, we

must supply a potential of just over +0.78 V to the reaction to make up for the

fact that E° is −0.78 V.

Cu(s) + Fe

(aq) n Fe(s) + Cu

(aq) E° =−0.78 V

When we do this, however, we note experimentally that the reaction still doesn’t

proceed toward products. If we supply a still greater positive potential, the reac-

tion does proceed. The extra voltage required is called the

overpotential of the

reaction. Overpotentials can be fairly high, especially when the products of

the reaction are gases. Unfortunately, we can’t easily predict what the overpo-

tential for a particular reaction will be. Instead we measure the overpotential

experimentally.

858 Chapter 19 Electrochemistry

gas

by-product

Exit port

Molten

/Na

AlF

mixture

Power source

Graphite

anodes

Molten

aluminum

FIGURE 19.19

The Hall–Heroult process. This process is still

used today to satisfy the world’s demand for

aluminum products. The product, molten alu-

minum, is relatively dense and settles to the

bottom of the electrochemical cell, where it is

drawn off and poured into castings.

The Applications of Electrolysis

The average American uses 142 tin cans each year. From

what are tin cans made? That might sound like a silly

question, but “tin” cans are actually made of steel

coated with a very thin layer of tin (Figure 19.20). Ap-

proximately 0.25% of the mass of a tin can is actually

tin, and chromium is becoming more common as a

coating on the steel. Without the coating, the steel

would rust and the contents would spoil. The tin coat-

ing on a steel can isn’t applied like paint on a house.

How is the tin applied to the steel?

The most common use for electrolysis is

electroplat-

ing

, or depositing metals onto a conducting surface.

The result is a coating that is very tightly integrated into

the surface of the metal. Because of this tight integra-

tion into the metal surface, the coating resists ﬂaking

and peeling. This coating makes the item more attrac-

tive and imparts some corrosion resistance or chemical

resistance to the surface. Electroplating is used in the manufacture of inexpensive

jewelry and chrome bumpers for your car, but most common electroplating

today occurs in the manufacture of tin cans.

Calculations Involving Electrolysis

The process used to coat a steel can with tin is an example of electrolysis. How is

it done? The steel can is hooked up to a power supply and dipped into a solution

of tin ions (Sn

). A block of tin metal is also placed in the solution and con-

nected to the power supply. Then a current is applied to the can (Figure 19.21). In

the terminology of the electrolytic cell, the block of tin metal becomes the anode

and the can becomes the site of reduction (the cathode) for tin ions.

Michael Faraday (remember him from our discussion on potentials and

spontaneity) noted that the amount of current applied to a cell is directly

proportional to the amount of metal that can be deposited in an electrolytic re-

action. We can represent this mathematically as

g =

A·s·(M

)



mol metal

mol e

−



where A is the number of amperes applied to the can

(1 ampere =1 coulomb of charge per second)

s is the number of seconds that the current is

applied

is the molar mass of the metal

F is Faraday’s constant (96,485 coulombs/mol

electrons)

the ratio (mol metal/mol e

−

) is the mole ratio

of the reduction half-reaction

It is helpful to look at this calculation as an extended

unit conversion problem. Exercise 19.12 shows how this

is done. Note that the unit conversion problem is the

same as the equation shown above.

19.7 Electrolytic Reactions 859

FIGURE 19.20

Tin cans are actually steel cans electroplated with a thin coat of tin. The

tin resists corrosion and helps keep the contents fresh.

Application

HEMICAL

ENCOUNTERS:

Electroplating

Power

source

Anode Cathode

–

Sn Sn

–

FIGURE 19.21

Electroplating a tin can. The tin can

acts as the cathode in the electrolysis

experiment.

Video Lesson: The Stoichiometry

of Electrolysis

EXERCISE 19.12 Golden Forks

To save money and still have a beautiful set of dinnerware, a chemist decides to elec-

troplate the metal with gold. How many grams of gold will be electroplated on a

fork if 2.5 A is applied to the fork for 20 s?

Solution

The number of grams can be calculated by starting with the current and perform-

ing a unit conversion. Note that the unit amperes (amps, A) can be written as

coulombs/second (C/s). We also need to examine the reduction half-reaction from

Table 19.3 to determine the number of electrons involved in the process.

(aq) + 3e

−

n Au(s)

2.5C

20 s

1 mol e

−

96,485 C

1 mol Au

3 mol e

−

196.97 g Au

1 mol Au

= 0.034 g Au

Amperes × time ×

faraday

× mole ratio × M

= grams

This means that our chemist would need to make sure to buy at least 0.034 g of gold

per fork.

PRACTICE 19.12

How many grams of tin will be deposited from a solution of tin(II) nitrate on a steel

can if 0.45 A are applied to the can for 1.5 h?

See Problems 73–78.

The calculations can also be done in reverse. If we want to coat our steel can

with a certain number of grams of tin, we can use either the equation or the unit

conversion method to calculate the number of amps that need to be applied.

As we’ve seen in this chapter, electrochemistry is a very useful topic, especially

in today’s society. Learning how muscles in our body begin a contraction can help

us understand how the heart works and why it is so important to maintain sufﬁ-

cient levels of electrolytes during physical exercise. And as we become increas-

ingly mobile, the demand for longer-lasting batteries will only increase. Batteries

power our cell phones, our portable CD players, and even our cars. Electroplating

the surfaces of many everyday items makes them appear expensive and confers

resistance to corrosion. In fact, everywhere we look, there’s electrochemistry!

860 Chapter 19 Electrochemistry

The Bottom Line

■

Redox reactions involve both a reduction and an oxi-

dation half-reaction. (Sections 19.1 and 19.3)

■

Redox reactions can be identiﬁed by determining the

oxidation state of the atoms involved in a reaction.

(Section 19.2)

■

Redox reactions can be balanced by summation of

balanced half-reactions. (Section 19.3)

■

Positive cell potentials indicate a spontaneous reac-

tion and are related to the free energy change by

∆G° =−nFE°. (Section 19.3)

■

Electrochemical cells require both a path for the

electrons and a path for other ions. (Section 19.4)

■

The oxidation reaction takes place at the anode. The

reduction reaction takes place at the cathode.

(Section 19.4)

■

The Nernst equation relates the actual potential of a

redox reaction to conditions other than the standard

conditions. (Section 19.6)

■

Half-reaction potentials can be used to determine

the relative reactivity of metals. Organization of the

metals in this fashion is known as a chemical reac-

tivity series. (Section 19.5)

■

Cell potentials enable us to calculate equilibrium

constants. (Section 19.6)

■

Electrowinning and electroplating are examples

of electrolysis reactions. In electrolysis, a positive

potential that includes the overpotential is applied

to force the reaction to run in reverse. (Section 19.7)

anode The electrode at which oxidation takes place.

(p. 826)

battery Two or more voltaic cells joined in series.

(p. 826)

cardiologist A heart specialist. (p. 824)

cathode The electrode at which reduction takes place.

(p. 826)

cell notation The chemist’s shorthand used to describe

electrochemical cells. (p. 847)

concentration cell A cell in which different concentra-

tions of identical ions on both sides of the cell

provide the driving force for the reaction. (p. 853)

corrosion The deterioration of a metal as a consequence

of oxidation. (p. 850)

disproportionation A reaction in which a single reactant

is both the oxidizing and the reducing agent.

(p. 830)

electrochemical cell A device that allows the exchange

between chemical and electrical energy. (p. 825)

electrochemistry The study of the reduction and oxida-

tion processes that occur at the meeting point of

different phases of a system. (p. 825)

electrochemists Scientists who create or analyze systems

that allow exchanges between chemical and electrical

energy. (p. 824)

electrode A metal surface that acts as a collector or

distributor for electrons. (p. 826)

electrodics The study of the interactions that occur

between a solution of electrolytes and an electrical

conductor, often a metal. (p. 825)

electrolytic cell A cell that requires the addition of

electrical energy to drive the chemical reaction under

study. (p. 826)

electromotive force (emf) A measure of how strongly a

species pulls electrons toward itself in a redox process.

Also known as voltage. (p. 834)

electroplating The process of depositing metals onto a

conducting surface. (p. 859)

electrowinning The isolation of pure metals from a

solution of metal ions. (p. 857)

Faraday’s constant A unit of electric charge equal to the

magnitude of charge on a mole of electrons. (p. 836)

fuel cell An electrochemical cell that utilizes continually

replaced oxidizing and reducing reagents to produce

electricity. (p. 826)

galvanic cell A cell that produces electricity from a

chemical reaction. Also known as a voltaic cell.

(p. 825)

gel electrophoresis The use of electric ﬁelds to separate

ions. (p. 825)

Key Words 861

Key Words

half-reaction An equation that describes the reduction

or oxidation part of a redox reaction. (p. 826)

Hall–Heroult process The most widely used process for

the preparation of aluminum from bauxite. (p. 858)

ionics The study of the behavior of ions dissolved in

liquids. (p. 825)

Nernst equation The equation used to determine the cell

potential at nonstandard conditions. (p. 852)

nitrogen cycle The path that nitrogen follows through its

different oxidation states on Earth. (p. 831)

overpotential The extra potential needed, above that

which is calculated, in order to make an electrochem-

ical process proceed. (p. 858)

oxidation The loss of electrons. (p. 826)

oxidation reaction In a redox reaction, the half-reaction

that supplies electrons. (p. 826)

oxidation state A bookkeeping tool that gives us insight

into the distribution of electrons in a compound.

Also known as oxidation number. (p. 827)

oxidizing agent A substance that causes the oxidation of

another substance. (p. 830)

potential The driving force (to perform a reaction) that

results from a difference in electrical charge between

two points. (p. 824)

reactivity series A ranking of the electrochemical reac-

tivity of some elements. (p. 849)

redox reactions Chemical reactions in which reduction

and oxidation occur. (p. 824)

reducing agent A substance that causes the reduction of

another substance. (p. 830)

reduction The gain of electrons. (p. 827)

reduction reaction In a redox reaction, the half-reaction

that acquires electrons. (p. 826)

sacriﬁcial anode A material that will oxidize more easily

than the one we seek to protect from oxidation.

(p. 850)

salt bridge A device containing a strong electrolyte

that allows ions to pass from beaker to beaker. (p. 845)

standard hydrogen electrode reaction (SHE) A reference

half-reaction of the reduction of hydrogen ion to

hydrogen gas, against which to compare our reduc-

tion. (p. 835)

standard potential (E°) The measure of the potential of a

reaction at standard conditions. (p. 834)

volt The SI unit of potential. (p. 834)

voltage A measure of how strongly a species pulls elec-

trons toward itself. Also known as electromotive force

(emf). (p. 834)

voltaic cell A cell that produces electricity from a chem-

ical reaction. Also known as a galvanic cell. (p. 825)

The answers to the odd-numbered problems and some selected

problems appear at the back of the book, as represented by the

blue numbering.

19.1 What Is Electrochemistry?

Skill Review

1. Explain why the descriptive term battery is often used, but is

technically not correct, to describe these most commonly

purchased sources of electricity.

2. Provide two other names for an electrochemical cell.

Chemical Applications and Practices

3. Every electrochemical cell is developed around two types of

chemical reactions. Name and describe both of these con-

nected reactions.

4. List three properties of electrochemical cells that are consid-

ered when designing an appropriate power source.

Section 19.2 Oxidation States—Electronic Bookkeeping

Skill Review

5. Determine the oxidation number of each atom in the struc-

ture of dimethylsulfoxide (DMSO).

10. Put the following compounds in order from the lowest to the

highest oxidation number for carbon.

OH CH

11. Determine the oxidation state for each atom in the following

compounds:

a. KMnO

b. LiMnO

c. NH

ClO

12. Determine the oxidation state for each atom in the following

compounds:

a. K

MnCl

b. Cr

c. C

13. Use the following four terms or expressions to identify each

of the chemical situations indicated: (is oxidized; is reduced; is

an oxidizing agent; is a reducing agent). You may use as many

terms as apply.

a. An atom has gained an electron.

b. An atom increases its oxidation number.

c. The oxidation number of an atom changes from −2 to –3.

14. Use the following four terms or expressions to identify each

of the chemical situations indicated: (is oxidized; is reduced; is

an oxidizing agent; is a reducing agent). You may use as many

terms as apply.

a. An atom decreases its oxidation number.

b. An atom loses two electrons.

c. The oxidation number of an atom changes from +3 to +5.

15. Supply the oxidation number of each atom, on both sides of

the reaction arrow, in the equation

3Mg(s) + 2H

(aq) n Mg

(PO

)

(aq) + 3H

(g)

16. Supply the oxidation number of each atom, on both sides of

the reaction arrow, in the equation

2AgNO

(aq) + Cu(s) n Cu(NO

)

(aq) + 2Ag(s)

Chemical Applications and Practices

17. The main active ingredient in commercial household bleach

is the hypochlorite ion, ClO

−

a. What is the oxidation number of the Cl atom in the ion?

b. Hypochlorite is known as a good oxidizing agent. Would

that property indicate that the Cl tends to gain or lose

electrons as it reacts? Explain the logic of your response.

18. The combustion of propane in a portable burner is a com-

mon redox reaction. Examine each component of the equa-

tion. Assign an oxidation number to each atom, and deter-

mine which component is acting as the reducing agent.

(g) + O

(g) n CO

(g) + H

O(g)

19. The brilliant red color of many ﬁreworks is due to the pres-

ence of strontium. However, strontium metal is not typically

found in its pure state. The following redox reaction depicts

the isolation of strontium from molten strontium chloride.

SrCl

(l) n Sr(s) + Cl

(g)

a. Assign an oxidation number to each of the atoms in the

reaction.

b. What has been reduced in the reaction?

862 Chapter 19 Electrochemistry

Focus Your Learning

6. Determine the oxidation number of each atom in the struc-

ture of periodic acid.

7. In which of the following compounds would the chlorine

atom have the most positive oxidation number? In which

would chlorine have the most negative oxidation number?

;ClO

;NaClO

;HCl

8. Which species in the following list shows the nitrogen atom

in its most reduced form? Which depicts nitrogen in its most

oxidized form? N

; HNO

;NH

9. Put the following compounds in order from the lowest to the

highest oxidation number for nitrogen.

NO N

ONO

20. Sodium thiosulfate is familiar to photographers as “hypo.” It

helps dissolve some of the silver salts used in developing pho-

tographs. Aqueous solutions of sodium thiosulfate can un-

dergo disproportionation reactions, as shown here:

2−

(aq) + H

(aq) n H

O(l) + S(s) + SO

(g)

a. What is the change in the oxidation number of S in S

2−

in the oxidation portion of the reaction?

b. What is the change in the oxidation number of S in S

2−

in the reduction portion of the reaction?

21. Four equations can be written to describe the rusting of iron:

(aq) + 2H

O(l) + 4e

−

n 4OH

−

(aq)

Fe(s) n Fe

(aq) + 2e

−

(aq) + 2OH

−

(aq) n Fe(OH)

(s)

Fe(OH)

(s) + OH

−

(aq) n FeO(OH)(s) + H

O(l) + e

−

a. Which of these are reduction or oxidation half-reactions?

b. What other kind of equation is present here?

c. Combine the equations to give the overall equation that

describes the formation of rust, FeO(OH)(s), from iron,

Fe(s).

22. In the chapter, we discussed the reaction of hydrogen and

oxygen to give water. How is it possible for this reaction to

proceed either as an explosion (as in the space shuttle main

engines) or as a gentle, readily managed source of electricity?

Section 19.3 Redox Equations

Skill Review

23. The superscript on the symbol E°

refers to standard condi-

tions for electrochemical reactions. What speciﬁcally does

that indicate for the concentrations and pressure of the react-

ing system?

24. The voltage values (emf) given on a standard reduction table

are referenced to a standard called SHE.

a. To what do the letters refer?

b. What is the voltage of the reference?

25. Complete each of these statements using the word positive,

negative, spontaneous,or nonspontaneous.

a. When E

° is negative, the electrochemical reaction is

___________.

b. When E° is positive, the value for



G° is ___________.

c. When a reaction is spontaneous, the values for



G° will

be __________, and the values for E° will be__________.

d. Nonspontaneous redox reactions have a ___________

value for E°.

26. The typical battery used in a standard ﬂashlight produces ap-

proximately 1.5 V. If the value for



G° of the reaction were

−289,500 J, what would you calculate as the moles of elec-

trons exchanged in the balanced redox reaction?

27. Combine these half-reactions in such a way that a galvanic

cell results, and calculate the cell potential.

+ e

−

n Fe

E° = 0.77 V

+ 2e

−

n Fe E° =−0.44 V

28. Combine these half-reactions in such a way that a galvanic

cell results, and calculate the cell potential.

+ 2e

−

n Sn E° =−0.14 V

+ 2e

−

n Sn

E° =+0.15 V

29. Calculate the free energy change for the cell in Problem 27.

30. Calculate the free energy change for the cell in Problem 28.

31. What is the free energy change for the following reaction at

standard conditions?

PbO

+4H

+2Hg +2Cl

−

n Pb

+2H

O +Hg

E°

cell

=1.12 V

32. What is the free energy change for the following reaction at

standard conditions?

+ 4H

+ 2Ni n 2H

O + 2Ni

E°

cell

= 1.46 V

33. What are three considerations or aspects that must be

“balanced” in a balanced redox reaction?

34. In this balanced redox reaction, chlorine is shown to replace

bromide ions from a solution.

(aq) + 2Br

−

(aq) n 2Cl

−

(aq) + Br

(aq)

What is the value of n in the overall reaction?

35. Balance these half-reactions in acidic solution. Is each a

reduction or an oxidation? Which substance is oxidized and

which is reduced?

a. CO

n H

b. Np

n NpO

36. Balance these half-reactions in acidic solution. Is each a re-

duction or an oxidation? Which substance is oxidized and

which is reduced?

a. I

n IO

−

b. NO

−

n NO

37. Balance these redox equations in acidic solution:

a. Sn

+ Cu

n Sn

+ Cu

b. S

2−

+ I

−

n S

2−

+ 3I

−

c. SO

−

+ Fe

n SO

2−

+ Fe

38. Balance these redox equations in acidic solution:

a. Al + Cu

n Al

+ Cu

b. UO

+Ag + Cl

−

n U

+AgCl

c. H

+ HBr n SO

+ Br

39. Balance the redox equation that illustrates the reaction of

solid copper and dichromate, ﬁrst in acidic and then in basic

solution.

Cu(s) + Cr

2−

(aq) n Cu

(aq) + Cr

(aq)

40. Balance the redox equation that illustrates the reaction of

permanganate and methanol, ﬁrst in acidic and then in basic

solution.

MnO

−

+ CH

OH n CO

2−

+ MnO

2−

41. Balance this redox reaction, ﬁrst in acidic and then in basic

solution.

ClO

−

+ I

−

n ClO

−

+ IO

−

42. Balance this redox reaction, ﬁrst in acidic and then in basic

solution.

Zn + NO

−

n Zn

+ NH

Focus Your Learning 863

Chemical Applications and Practices

43. Using the standard reduction potentials found in the appen-

dix, locate the half-cell reaction for zinc. Zinc is often used in

the production of dry cell batteries. It is also used to protect

other metals from oxidation.

a. What is the E

° value for this reaction?

b. What would be the value of



° for the half-reaction?

c. What would be the E

° value if the reaction stoichiometry

were doubled?

44. Splitting water into hydrogen gas and oxygen gas is one tech-

nique being investigated as a way to produce hydrogen gas for

use in fuel cells. The reaction is shown here as

O n 2H

+ O

If the E° value for this nonspontaneous reaction were ap-

proximately −2.00 V, what would you calculate as the value

for



°?

45. Use the Standard Table of Reduction Potentials to answer the

following questions:

a. Which of the three alkali metal ions Na, K, and Li has the

least potential to attract electrons?

b. Of the three halogens which would be the strongest oxi-

dizing agent, F

,Cl

,or Br

c. If you prepared a battery by connecting two of the follow-

ing three half-cells, which combination could produce the

highest potential? Zn, Cu, Ni

46. Suppose you attempted to build a battery using lead (and

) with chromium (and Cr

a. Write out the half-cell reduction reactions and their E

values for each.

b. Which of the two would provide the reduction reaction,

and which would provide the oxidation reaction?

c. What would be the total E

of the overall redox reaction?

d. How many moles of electrons would be exchanged in the

overall reaction?

47. The oxidation of copper metal gives rise to a beautiful green

patina. An equation that illustrates the oxidation of copper is

shown here. Balance this redox reaction.

Cu(s) + CO

(g) n CuO(s) + C

2−

(aq)

48. One technique used for the detection of ethanol involves the

following redox reaction with dichromate ions. Balance this

redox reaction.

OH(aq) + Cr

2−

(aq) n CH

H(aq) + Cr

(aq)

49. As with acid–base reactions, careful addition of a solution

containing an oxidizing agent to a solution containing a re-

ducing agent can be used for titration analysis. Standardizing

a solution of potassium permanganate to be used later in a

redox titration often involves reaction with a known amount

of sodium oxalate. Balance the redox reaction used in the

standardization process.

2−

(aq) + MnO

−

(aq) n CO

(g) + Mn

(aq) + H

O(l)

50. Once a standard solution of potassium permanganate is pre-

pared, it can be used to determine the concentration of iron

in an unknown sample. For example, the iron content of a

small steel sample could be obtained through a titration re-

action with a standard permanganate solution. The redox

reaction that would take place in the analysis is shown here

(unbalanced, and after the iron has been prepared as +2 ion).

Balance the redox titration and identify the reducing agent.

MnO

−

(aq) + Fe

(aq) n Mn

(aq) + Fe

(aq) + H

O(l)

51. One step in a common method for the analysis of vitamin C

in juice drinks is to oxidize the vitamin C (ascorbic acid) with

aqueous I

. Balance this redox reaction.

(aq) + I

(aq) n I

−

(aq) + C

(aq)

52. There are many half-cell reaction combinations that could be

used to make batteries. One such reaction would be to use

system consisting of silver(II) oxide and zinc. The reaction

shown below would have to occur in a basic environment.

Balance this redox reaction.

Zn(s) +AgO(s) n Zn(OH)

(s) +Ag(s)

Section 19.4 Electrochemical Cells

Skill Review

53. Deﬁne the following terms in your own words: cell, half-

reaction, galvanic cell, voltaic cell, electromotive force.

54. Compare and contrast an electrolytic cell with a galvanic cell

on the basis of sign on the E° value, ability to do work, chem-

ical process taking place at the anode, and spontaneity.

55. The following notation refers to a speciﬁc voltaic cell. Iden-

tify the species that would serve as the anode and the species

that is the oxidizing agent.

Zn(s) | Zn

(aq) || Cu

(aq) | Cu(s)

56. The following notation refers to a speciﬁc voltaic cell. Iden-

tify the species that would serve as the anode and the species

that is the oxidizing agent.

Mg(s) | Mg

(aq) || Co

(aq) | Co(s)

57. The following schematic diagram represents two metals and

their cations in separate beakers joined by a NaNO

salt

bridge. If the metals were lead and silver (with their respec-

tive cations Pb

and Ag

), which beaker would require the

silver and which the lead if the nitrate ions were moving from

left to right through the salt bridge? Justify your answer.

864 Chapter 19 Electrochemistry

58. The following schematic diagram represents two metals and

their cations in separate beakers joined by a NaNO

salt

bridge. If the metals were copper and iron (with their respec-

tive cations Cu

and Fe

), which beaker would require the

copper and which the iron if the nitrate ions were moving

from left to right through the salt bridge? Justify your answer.

c. In which direction, to the left or to the right, would the

following reaction take place spontaneously? (Assume

standard conditions.)

Ag + Fe

n Fe

+Ag

64. Under standard conditions one of the following two reac-

tions will not occur spontaneously. Use the Standard Table of

Reduction Potentials to explain how you could correctly pre-

dict which one will not take place.

Fe + Sn

n Fe

+ Sn

+ Fe n Fe

+ Cu

Chemical Applications and Practices

65. Suppose you have a shiny piece of metal that is unlabeled. It

is known to be either aluminum or tin. You also have a solu-

tion of 1.0 M nickel(II) nitrate. Suggest a chemical test that

you could use to determine the identity of the metal using the

solution. Explain your expected results and the basis of your

conclusion.

66. One technique to protect against oxidation of structures such

as ship hulls and underground iron pipes is to place the

structure in contact with a metal that will oxidize more easily

than the iron in the structure. The more active metal is then

referred to as a sacriﬁcial anode. Neglecting such issues as

cost and availability, would copper or zinc make the better

sacriﬁcial anode in an attempt to protect an iron-based struc-

ture? Explain the basis of your choice.

Section 19.6 Not-So-Standard Conditions:

The Nernst Equation

Skill Review

67. The constant used in the Nernst equation, 0.0257, is derived

from the combination of three other constants. Obtain the

three values, assuming standard temperature, and derive the

constant used in the equation. Some texts refer to the Nernst

equation using the base 10 logarithmic scale. What would be

the value of the constant in that scale?

68. The simpliﬁed equation presented here represents the redox

reaction taking place inside the typical, nonalkaline ﬂashlight

battery. Use the principles described in the Nernst equation

to answer the questions that follow.

Zn(s) + 2MnO

(s) + 2NH

(aq) n

(aq) + Mn

(s) + 2NH

(g) + H

O(l)

a. What would be the effect on the spontaneity of the re-

action if the concentration of NH

(aq) were decreased?

b. What would be the effect on the voltage of the cell if the

concentration of NH

(aq) were decreased?

c. What is the value of n in the equation?

d. At equilibrium what is the value of E?

69. Calculate the value of E

cell

for the reaction of iron and cop-

per(II), given the speciﬁc concentrations listed.

a. Fe(s) | Fe

(0.10 M) || Cu

(0.10 M) | Cu(s)

b. Fe(s) | Fe

(1.5 M) || Cu

(0.10 M) | Cu(s)

c. Fe(s) | Fe

(0.10 M) || Cu

(1.5 M) | Cu(s)

70. Calculate the value of E

cell

for the reaction of cobalt(II) chlo-

ride and zinc metal, given the speciﬁc concentrations listed.

a. Zn(s) | Zn

(0.10 M) || Co

(0.10 M) | Co(s)

b. Zn(s) | Zn

(2.5 M) || Co

(0.050 M) | Co(s)

c. Zn(s) | Zn

(0.010 M) || Co

(0.10 M) | Co(s)

Focus Your Learning

865

Chemical Applications and Practices

59. Diagram a battery consisting of two beakers and a salt bridge

that makes use of the two half-reactions silver and gold (sort

of expensive, but you’re worth it).

a. Predict the E

° value for the battery.

b. What is the value of n for the balanced equation?

c. Label the cathode.

d. Exactly what species is acting as the oxidizing agent?

e. Represent the battery in the ABC notation.

60. Search Internet references using the keywords “hydrogen fuel

cell” and report on the basic operation of such a cell. Give

particular emphasis to the electrolyte used in such a cell. Be

sure to report the necessary URL for your references.

61. Small button-sized batteries can be made using mercury and

zinc. If the two half-reactions are as follows, what would you

report as the ABC notation for the battery? (Mercury in the

batteries is considered a toxic substance and should be care-

fully recycled.)

Zn(s) n ZnO(s)

HgO(s) n Hg(l)

Note that the reactions are not balanced. The reaction takes

place in a basic medium.

62. Write out the standard reduction half-cell reactions and E

values for Zn, Ni, Pb, and H

(g).

a. Select the two half-reactions that, when combined, would

produce the battery with the greatest theoretical overall E

value, and report that value.

b. Which, if any, of those represented would have electrons

move along an external wire, when connected with the

hydrogen half-cell, toward the hydrogen electrode?

Section 19.5 Chemical Reactivity Series

Skill Review

63. Use the Standard Table of Reduction Potentials to answer the

following questions.

a. Which is the better reducing agent, Ba or Ca?

b. Which is the better oxidizing agent, Pb

or Ni

Chemical Applications and Practices

71. Suppose that at night on a campout your ﬂashlight dims and

ﬁnally stops working. Your friend remarks, “I guess the bat-

tery is dead.” In what three other chemical ways, using Gibbs

free energy, equilibrium, and potential, could you state the

same conclusion as your friend?

72. A quick source of hydrogen gas in the lab is to carefully place

some solid magnesium in hydrochloric acid.

Mg(s) + 2HCl(aq) n MgCl

(aq) + H

(g)

If the Mg

concentration were 1.0 M and the HCl concen-

tration were 0.10 M (assume completely dissociated HCl),

what would you calculate as the E value for the reaction at

25°C?

Mg(s) | Mg

(aq) || H

(0.1 M) | H

(1 atm)

Section 19.7 Electrolytic Reactions

Skill Review

73. Calculate the number of grams of gold electroplated

onto a surface given these conditions. (Assume that the solu-

tion of Au

ions is concentrated enough to complete the

electrolysis.)

a. 1.25 A for 60 s b. 2.11 A for 2.33 h c. 0.75 A for 1 d

74. Calculate the time required to electroplate 1.0 g of tin metal

onto a surface given these conditions. (Assume that the solu-

tion of Sn

ions is concentrated enough to complete the

electrolysis.)

a. 2.25 A b. 0.11 A c. 1.38 A

75. Alessandro Volta is given historical precedence in the discov-

ery of the ﬁrst battery in the sense that we think of batteries

today. The “Voltaic pile” consisted of dissimilar metals joined

by salt-moistened paper strips. Davy and his assistant

Michael Faraday later reﬁned this. At that time, early electro-

plating businesses began developing in England. If such a

business silver-plated a teaspoon, would the teaspoon be the

cathode or the anode in such a process? If the spoon received

0.33 g of silver after being plated for 1.0 h, what amperage

was used?

76. In addition to isolating sodium, Sir Humphrey Davy isolated

potassium, calcium, and magnesium from impure natural

ore samples.

a. If he used a current of 1.00 A for 1.00 h, how many grams

of magnesium metal could he obtain from molten MgCl

b. Using the same electrical set up, how many grams of

potassium metal could Sir Humphrey isolate from molten

KCl?

Chemical Applications and Practices

77. In the typical lead storage battery found in most automo-

biles, lead is oxidized to Pb

(in the form of PbSO

). In

recharging, the reaction is reversed. If a battery were

recharged for 30.0 min at a current of 8.00 A, how many

grams of lead would be reduced, from PbSO

during the

process?

78. Germanium has become a valuable metal in semiconductor

ﬁelds. If a 1.00 A current were used for 1.00 h to plate out

0.677 g of germanium, what would you calculate as the

oxidation number of the germanium ions in the plating

solution?

Comprehensive Problems

79. Use the illustration below as the starting point to draw the

electrolytic cell that would be used to plate copper metal onto

a steel saucepan. Be sure to indicate the location of the cath-

ode and the anode, the location of the copper electrode, and

the steel saucepan.

866 Chapter 19 Electrochemistry

Power

source

Anode Cathode

80. The Nernst equation can be used just as well with half-cell re-

actions as with balanced redox reactions. An application of

this is the use of potential differences in the acid level of so-

lutions compared to a standard when producing the probes

for pH meters. Using the half-reaction 2H

(aq) + 2e

−

(g) at 1 atm, what would you calculate, using the Nernst

equation, as the value for E in these situations?

a. Pure water, pH = 7.00

b. An acid solution with pH = 2.00

c. A 0.10 M solution of nitric acid

d. A 0.10 M solution of acetic acid, K

= 1.8

−5

81. Potassium ferrate (K

FeO

) is a powerful oxidizing agent

° = 2.20 V relative to SHE in acidic solution) in which the

iron(VI) ion is reduced to the iron(III) ion. It oxidizes water

to oxygen gas.

a. Determine the cell potential and equilibrium constant for

the reaction at standard conditions for the oxidation of

water by ferrate ion in acidic solution.

b. Balance the reaction in basic solution.

c. Calculate the cell potential, given the following initial

concentrations:

[FeO

2−

] = 1.5

−3

M;[Fe

] = 1.1

−3

= 8.3

−5

atm; pH = 2.8

d. The reaction results in the production of a yellow solid

that is especially apparent at high pH. Can you suggest

what this might be? Further, can you suggest how this solid

might help if the ferrate ion were used to treat wastewater

contamination?

82. Potassium permanganate (KMnO

) is a useful analytical

reagent for determining the percentage of iron in an iron ore.

The procedure includes dissolving the iron with HCl and

then converting all of the iron in the ore to Fe

using several

reagents. The titration of the resulting Fe

with MnO

−

+ MnO

−

n Fe

+ Mn

A sample of the original iron ore weighing 3.852 g was

processed for titration in a total volume of 300.0 mL of solu-

tion. A 100.0 mL aliquot (accurately measured portion) was

removed and titrated with a 0.1025 M KMnO

solution. A

total of 23.14 mL was required to the light pink endpoint.

What is the percent of iron in the ore sample?

83. We noted in Chapter 14, Thermodynamics, that we can have

“standard conditions” at a given pressure, but the tempera-

ture can vary and must be stated as part of specifying your

standard values. If we deﬁne our standard values at 1 bar

pressure and 3000°C, how would that affect the values for our

“new” standard reduction potentials compared to the ones

we commonly use at 25°C?

84. We discuss rechargeable batteries in Section 19.4. You know

from your own experience that these batteries can be

recharged hundreds of times but eventually wear down and

must be replaced. Why can’t they be recharged an inﬁnite

number of times? In particular, investigate the way in which

metals plate back onto the cathode when the recharging takes

place.

85. Chrome plating is the process by which Cr

is electroplated

on automobile surfaces to give them a nice luster. The

chromium layer can be as little as 10.0 µm thick. How many

grams of CrO

would be necessary to plate the back surface

of a car mirror that has a coverable area of 150 cm

(d = 7.2 g/cm

86. We discuss fuel cells at several places in the chapter, but never

actually calculate the electrochemical data for it. Calculate

the E°, G° and K at 298K for the hydrogen/oxygen fuel cell

described in Section 19.2. Once you have done so, can you

propose a fuel cell that might be more effective for use in au-

tomobiles? What are the criteria that you use to make that

judgment?

Thinking Beyond the Calculation

87. A battery designer wishes to prepare a solution-based battery

for use in a new automobile. The designer chooses iron and

zinc as the two metals for study.

a. Draw the setup (using beakers and a salt bridge) that

indicates the location of the iron and zinc electrodes, the

iron(II) and zinc(II) solutions, and the ﬂow of electrons.

b. Which half-reaction is the oxidation reaction?

c. What is the cell potential for the reaction if the initial

concentrations of iron(II) and zinc(II) are 0.25 M at 25°C?

d. What would be the cell potential if the concentrations of

both solutions in part c were increased to 1.0 M? ...to

2.0 M?

e. If the battery is cooled to 10°C, is there a change in the cell

potential? If so, what is the new potential?

f. Describe at least one advantage to using a battery made

from these metals.

g. Describe at least one disadvantage to the use of iron in a

battery.

Focus Your Learning

867