Electron Transfer Reactions 831 Chapter 17 Test Bank Docx

CHAPTER 17

ELECTRON TRANSFER REACTIONS

CHAPTER STUDY OBJECTIVES

1. Use oxidation numbers to show what is being oxidized and what is being reduced in a redox reaction.

SKILLS TO MASTER: Recognizing the oxidant and the reducing agent; assigning oxidation numbers

KEY CONCEPTS: Oxidation is the loss of electrons from a substance. Reduction is the gain of electrons by a substance. Oxidation and reduction always occur together.

2. Balance redox reactions using the half-reaction method.

SKILLS TO MASTER: Identifying half-reactions; balancing half-reactions; combining half-reactions

KEY CONCEPTS: Redox reactions can be separated into two half-reactions, one for the oxidation and one for the reduction.

3. Describe galvanic cells.

SKILLS TO MASTER: Describing a galvanic cell from the balanced redox reaction; writing the shorthand notation of an electrochemical cell from the balanced redox reaction and vice-versa

KEY CONCEPTS: Oxidation always occurs at the anode and reduction always occurs at the cathode. A galvanic cell has a spontaneous redox reaction.

4. Calculate standard cell potentials.

SKILLS TO MASTER: Calculating standard cell potentials from standard reduction potentials

KEY CONCEPTS: All reduction potentials are relative to one another. The standard potential for reducing H₃O⁺(aq) to H₂(g) and H₂O(l) is defined to be 0.00 V. In a galvanic cell, the half-reaction with the more negative reduction potential occurs at the anode as an oxidation. The half-reaction with the more positive reduction potential occurs at the cathode as a reduction. A galvanic cell has a positive standard cell potential. When a half-reaction is multiplied by any integer, its potential remains unchanged.

5. Relate cell potential to the reaction conditions.

SKILLS TO MASTER: Calculating the free energy change from the cell potential; calculating equilibrium constants of redox reactions; calculating the cell potential under non-standard

Conditions using the Nernst equation; performing stoichiometric electrochemical calculations

KEY CONCEPTS: A cell potential depends on the concentrations of all species in the reaction quotient and on the temperature.

6. Explain the chemistry of everyday redox reactions.

SKILLS TO MASTER: Explaining the chemistries of several types of batteries; explaining under what conditions corrosion happens

KEY CONCEPTS: Batteries are galvanic cells. A fuel cell is a galvanic cell requiring a continuous supply of fuel and oxidant.

7. Explain electrolytic reactions and cells.

SKILLS TO MASTER: Predicting the reactions that occur at each electrode in an electrolytic cell; doing quantitative electroplating calculations

KEY CONCEPTS: An electrolytic cell has a non-spontaneous redox reaction and a negative standard cell potential.

Multiple Choice QUESTIONS

1. Which of the following is NOT a redox reaction?

a) 4 NH₃ + 5 O₂ 🡪 4 NO + 6 H₂O

b) 2 CO + O₂ 🡪 2 CO₂

c) S + 2 F₂🡪 SF₄

d) AgNO₃ + KI 🡪 AgI(s) + KNO₃

e) Cl₂ + 2H₂O 🡪 2Cl^- + 2OCl^- + 4H⁺

Difficulty: Medium

Learning Objective: Use oxidation numbers to show what is being oxidized and what is being reduced in a redox reaction.

Section Reference: 17.1 Recognizing Redox Reactions

2. Which of the following are redox reactions?

I. 2 N₂ + 3 H₂ 🡪 2 NH₃

II. 4 Al + 3 O₂ 🡪 2 Al₂O₃

III. 2 NO₂ 🡪 N₂O₄

IV. FeCl₃ (aq) + 6 NH₃(aq) 🡪 Fe(NH₃)₆³⁺ (aq) + 3 Cl^- (aq)

V. Cu²⁺ + Zn 🡪 Cu + Zn²⁺

a) I, IV and V

b) All but III

c) I and II

d) I, II and IV

e) I, II and V

Difficulty: Medium

Learning Objective: Use oxidation numbers to show what is being oxidized and what is being reduced in a redox reaction.

Section Reference: 17.1 Recognizing Redox Reactions

3. Which elements are changing oxidation states in the following reaction?

Zn(s) + 2 MnO₂(s) + H₂O(l) 🡪 Zn(OH)₂(s) + Mn₂O₃(s)

a) Zn, O

b) Mn, Zn

c) H, Mn

d) O, Mn

e) H, O

Difficulty: Easy

Learning Objective: Use oxidation numbers to show what is being oxidized and what is being reduced in a redox reaction.

Section Reference: 17.1 Recognizing Redox Reactions

4. Assign oxidation numbers to all the elements in LiAlH₄.

a) Li, +1; Al, +3; H, +1

b) Li, +1; Al, –5; H, +1

c) Li, +1; Al, +3; H, –1

d) Li, +1; Al, –2; H, +1

e) Li, +1; Al, 0; H, –1

Difficulty: Medium

Learning Objective: Use oxidation numbers to show what is being oxidized and what is being reduced in a redox reaction.

Section Reference: 17.1 Recognizing Redox Reactions

Feedback: a) common oxidation states all elements; b) common oxidation states of Li and H; c) correct answer; d) common oxidation states of Li and H; e) failure to count all four hydrogens

5. Assign oxidation numbers to all the elements in KCN.

a) K, +1; C +4; N, –3

b) K, +1; C +2; N, –3

c) K, +1; C –4; N, +3

d) K, +1; C –6; N, +5

e) K, +1; C,+3; N –4

Difficulty: Easy

Learning Objective: Use oxidation numbers to show what is being oxidized and what is being reduced in a redox reaction.

Section Reference: 17.1 Recognizing Redox Reactions

Feedback: a) common oxidation states all elements; b) correct answer; c) oxidation state on nitrogen wrong; d) common oxidation states for K and N

6. Nitrogen has many possible oxidation numbers; put the following nitrogen compounds in order of increasing oxidation number: NO₂, HNO₃, NO₂^-, NO.

a) NO₂, HNO₃, NO₂^-, NO

b) NO, HNO₃, NO₂^-, NO₂

c) HNO₃, NO₂, NO₂^-, NO

d) NO, NO₂^-, NO₂, HNO₃

e) NO₂^-, NO, NO₂, HNO₃

Difficulty: Easy

Learning Objective: Use oxidation numbers to show what is being oxidized and what is being reduced in a redox reaction.

Section Reference: 17.1 Recognizing Redox Reactions

7. Determine the coefficient for Sn⁺² in the following balanced redox reaction.

MnO₄^- + Sn²⁺ 🡪 Sn⁴⁺ + Mn²⁺ (acidic solution)

a) 2

b) 4

c) 5

d) 6

e) 10

Difficulty: Medium

Learning Objective: Balance redox reactions using the half-reaction method.

Section Reference: 17.2 Balancing Redox Reactions

8. Consider the redox reaction of permanganate and sulphur:

MnO₄^- + S 🡪 Mn²⁺ + SO₄^2- (acidic solution)

If the coefficient of MnO₄^- is 6 in the balanced equation, what is the coefficient of H₂O?

a) 2

b) 3

c) 4

d) 5

e) 6

Difficulty: Medium

Learning Objective: Balance redox reactions using the half-reaction method.

Section Reference: 17.2 Balancing Redox Reactions

9. Consider the redox reaction of nitric acid and copper:

Cu + HNO₃ 🡪 Cu(NO₃)₂ + NO (acidic solution)

If the coefficient of Cu is 3 in the balanced equation, what is the coefficient of HNO₃?

a) 4

b) 5

c) 7

d) 8

e) 10

Difficulty: Medium

Learning Objective: Balance redox reactions using the half-reaction method.

Section Reference: 17.2 Balancing Redox Reactions

10. Consider the redox reaction of triiodide and oxygen:

I₃^- + O₂ 🡪 I₂ + OH^- (basic solution)

If the coefficient of I₃^- is 4 in the balanced equation, what is the coefficient of OH^-?

a) 2

b) 4

c) 6

d) 8

e) 10

Difficulty: Medium

Learning Objective: Balance redox reactions using the half-reaction method.

Section Reference: 17.2 Balancing Redox Reactions

11. The following reaction occurs in a galvanic cell:

NiO₂ + Cd + H₂O 🡪 Cd(OH)₂ + Ni(OH)₂ + 2 OH^-

Which redox process in this battery occurs at a passive electrode?

a) Cd 🡪 Cd(OH)₂

b) NiO₂ 🡪 Ni(OH)₂

c) O₂ 🡪 4 OH^-

d) H₂O 🡪 OH^-

e) neither electrode is passive

Difficulty: Medium

Learning Objective: Describe galvanic cells.

Section Reference: 17.3 Galvanic Cells

12. What is the role of the electrolyte in a galvanic cell?

a) to facilitate rapid diffusion of the redox reagents to each other

b) to facilitate electron transport though the solution

c) to complete the electrical circuit by ion transport

d) to supply the ions for precipitating redox products

e) to protect electrodes from corrosion

Difficulty: Easy

Learning Objective: Describe galvanic cells.

Section Reference: 17.3 Galvanic Cells

13. For the reaction given, which half reaction occurs at the cathode?

NiO₂ + Cd + H₂O 🡪 Cd(OH)₂ + Ni(OH)₂ + 2 OH^-

a) Cd 🡪 Cd(OH)₂

b) NiO₂ 🡪 Ni(OH)₂

c) Cd(OH)₂ 🡪 Cd

d) H₂O 🡪 OH^-

e) There is no cathode as the cathodic reaction always occurs at the passive electrode.

Difficulty: Medium

Learning Objective: Describe galvanic cells.

Section Reference: 17.3 Galvanic Cells

14. For the reaction given below, which half reaction occurs at the cathode?

Pb(s) + PbO₂(s) + 2HSO₄^-(aq) + 2H₃O⁺(aq) 🡪 2PbSO₄(aq) +4H₂O(l)

a) Pb 🡪 PbSO₄

b) PbO₂ 🡪 PbSO₄

c) HSO₄^- 🡪 PbSO₄

d) H₃O⁺ 🡪 H₂O

e) PbSO₄ 🡪 PbO₂

Difficulty: Medium

Learning Objective: Describe galvanic cells.

Section Reference: 17.3 Galvanic Cells

Feedback: a) is oxidation reaction at anode; b) correct answer; c) neither oxidation or reduction; d) this is not a redox reaction; c) this is an oxidation reaction, but does not occur in the spontaneous cell

15. For the reaction given below, identify the anode and describe what happens to the electrode as the reaction continues.

3Fe(s) + Cr₂O₇^2-(aq) + 14 H⁺(aq) 🡪 3Fe²⁺(aq) + 2 Cr⁺³(aq) + 7H₂O(l)

a) Fe, converted to Fe²⁺; electrode decreases in size

b) Fe, converted to Fe²⁺; electrode gains mass

c) Passive electrode at which Cr₂O₇^2-(aq) is oxidized; electrode is unchanged

d) Fe electrode at which Cr⁺³ is oxidized; electrode is unchanged

e) Fe electrode; no change

Difficulty: Medium

Learning Objective: Describe galvanic cells.

Section Reference: 17.3 Galvanic Cells

Feedback: a) correct answer; b) anode correctly identified, unclear on the process occurring; c) Cr₂O₇^2- is reduced; student unclear on meaning of oxidation; d) incorrect identification of oxidation process at anode; e) correctly identified oxidation process, unclear on physical changes occurring

16. For the reaction given below, which half reaction occurs at the anode?

2 H₂(g) + O₂(g) 🡪 2 H₂O(l)

a) H₂ 🡪 H₂O

b) O₂ 🡪 H₂O

c) H₂O 🡪 H₂

d) H₂O 🡪 O₂

e) H₂(g) 🡪 2H⁺

Difficulty: Medium

Learning Objective: Describe galvanic cells.

Section Reference: 17.3 Galvanic Cells

17. Consult a table of reduction potentials (Table 17-1) and determine which two metals are capable of reducing iron (II) to iron under standard conditions.

a) Ca, Sn

b) Sn, Pb

c) Al, Mg

d) Mg, Cu

e) Cd, Hg

Difficulty: Easy

Learning Objective: Calculate standard cell potentials.

Section Reference: 17.4 Cell Potentials

18. Platinum metal is quite resistant to oxidation as may be deduced by its reduction potential:

Pt²⁺ + 2e 🡪 Pt E° ≈ 1.2 V

Examine a table of reduction potentials (Table 17-1) and determine two elements capable of oxidizing platinum under standard conditions.

a) Au, F₂

b) F₂, Fe

c) F₂, Cl₂

d) Br₂, Ag

e) Mn, Au

Difficulty: Easy

Learning Objective: Calculate standard cell potentials.

Section Reference: 17.4 Cell Potentials

19. Consider an electrochemical cell consisting of an Fe(s) electrode, Fe(NO₃)₂ electrolyte connected through a salt bridge to a Ag wire coated in AgCl(s) immersed in an aqueous KCl solution. Is the standard cell and balanced galvanic cell reaction:

(If needed, refer to Table 17-1.)

a) 0.259 V, Fe(s) + 3AgCl(s) 🡪 Fe³⁺(aq) + 3Ag(s) + 3Cl^-(aq)

b) 0.669 V, Fe(s) + 2AgCl(s) 🡪 Fe²⁺(aq) + 2Ag(s) + 2Cl^-(aq)

c) –0.669 V, Fe(s) + 2AgCl(s) 🡪 Fe²⁺(aq) + 3Ag(s) + 3Cl^-(aq)

d) –0.225 V, Fe(s) + 2AgCl(s) 🡪 Fe²⁺(aq) + 2Ag(s) + 2Cl^-(aq)

e) 0.669 V, Fe(s) + AgCl(s) 🡪 Fe²⁺(aq) + Ag(s) + Cl^-(aq)

Difficulty: Hard

Learning Objective: Calculate standard cell potentials.

Section Reference: 17.4 Cell Potentials

Feedback: a) incorrect oxidation reaction; b) correct answer; c) sign reversed on cell potential; d) cell potential calculated as E^o_cathode+ E^o_anode; e) reaction not balanced

20. Consider the Daniell cell where the cell reaction and standard potential are:

Zn(s) + Cu²⁺ (aq) 🡪 Zn²⁺ (aq) + Cu (s) E° = 1.10 V

If the cell is initially at standard conditions ([Cu²⁺] = [Zn²⁺] = 1.00 M), what are the concentrations of Cu²⁺ and Zn²⁺ when the cell potential has fallen to 1.06 V?

If needed, use the following equation:

ΔG˚ = –nFE˚, E˚ , E = E˚ – , moles e^- =

(If needed, refer to Table 17-1.)

a) [Cu²⁺]= 8.5 x 10^-2 M; [Zn²⁺]=1.91 M

b) [Cu²⁺]= 0.94 M; [Zn²⁺]= 1.06 M

c) [Cu²⁺]= 1.91 M; [Zn²⁺]=8.5 x 10^-2 M

d) [Cu²⁺]= 0.50 M; [Zn²⁺]=1.50 M

e) [Cu²⁺]= 0.90 M; [Zn²⁺]=1.10 M

Difficulty: Hard

Learning Objective: Relate cell potential to the reaction conditions.

Section Reference: 17.5 Free Energy and Electrochemistry

21. What is the correct description, in line notation, for an electrochemical cell comprised of only Ag wire, AgNO₃ electrolyte solution, and a salt bridge having ΔG = –2 kJ?

If needed, use the following equation:

ΔG˚ = –nFE˚, E˚ , E = E˚ – , moles e^- =

(If needed, refer to Table 17-1.)

a) Ag(s) ⎜Ag⁺( aq, 1.00 M) ⎢⎢Ag⁺(aq, 2.25 M) ⎢Ag(s)

b) Ag(s) ⎜Ag⁺( aq, 1.00 M) ⎢⎢Ag⁺(aq, 1.00 M) ⎢Ag(s)

c) Ag(s) ⎜Ag⁺( aq, 2.25 M) ⎢⎢Ag⁺(aq, 1.00 M) ⎢Ag(s)

d) Ag(s) ⎜Ag⁺( aq, 0.445 M) ⎢⎢Ag⁺(aq, 1.00 M) ⎢Ag(s)

e) Ag(s) ⎜Ag⁺( aq, 1.00 M), Ag⁺(aq, 2.25 M) ⎢Ag(s)

Difficulty: Hard

Learning Objective: Relate cell potential to the reaction conditions.

Section Reference: 17.5 Free Energy and Electrochemistry

Feedback: a) concentrations reversed; b) E = 0 V,; c) correct answer; d) concentrations reversed; e) incorrect notation

22. For the working galvanic cell shown at standard conditions, how would you increase the cell potential?

(If needed, refer to Table 17-1.)

a) Make the Pt electrode larger.

b) Make the Copper electrode larger.

c) Increase the concentration of KI.

d) Increase the concentration of I₂.

e) Make the Cu electrode smaller.

Difficulty: Medium

Learning Objective: Relate cell potential to the reaction conditions.

Section Reference: 17.5 Free Energy and Electrochemistry

23. For the following galvanic cell what will be its potential when the reaction reaches equilibrium?

(If needed, refer to Table 17-1.)

a) 0.0 V

b) 0.458 V

c) 1.142 V

d) 0.272 V

e) 1.26 V

Difficulty: Medium

Learning Objective: Relate cell potential to the reaction conditions.

Section Reference: 17.5 Free Energy and Electrochemistry

24. Ships, storage tanks, and other large metal items may be protected from corrosion by

(If needed, refer to Table 17-1.)

a) a sacrificial cathode.

b) reduction of K⁺ to K.

c) a sacrificial anode.

d) coating with potassium metal.

e) there is no way to protect metals from corrosion.

Difficulty: Easy

Learning Objective: Explain the chemistry of everyday redox reactions.

Section Reference: 17.6 Redox in Action

25. What are the possible oxidation states of corroded iron?

(If needed, refer to Table 17-1.)

a) 2

b) 0

c) 3

d) 2 and 3

e) 0, 2 and 3

Difficulty: Easy

Learning Objective: Explain the chemistry of everyday redox reactions.

Section Reference: 17.6 Redox in Action

Feedback: a) 1 of 2 possible oxidation states; b) no corrosion has occurred; c) 1 of 2 possible oxidation states; d) correct answer; e) iron in oxidation state 0 has not undergone corrosion

26. How is aluminium protected from oxidation?

(If needed, refer to Table 17-1.)

a) formation of an oxide layer

b) attaching it to a block of zinc

c) coating it with iron

d) Aluminium is not protected from oxidation, since E^o for Al is –1.662 V it will always corrode.

e) Provide a protective paint on the surface of the aluminium metal.

Difficulty: Easy

Learning Objective: Explain the chemistry of everyday redox reactions.

Section Reference: 17.6 Redox in Action

27. Which of the following combinations would provide the largest potential for a battery?

(If needed, refer to Table 17-1.)

a) Br₂ and Fe

b) Br^-1 and Fe⁺²

c) Al and Cu⁺²

d) Al⁺³ and Cu⁺²

e) Al and Br^-

Difficulty: Medium

Learning Objective: Explain the chemistry of everyday redox reactions.

Section Reference: 17.6 Redox in Action

28. You have an abundant supply of NaCl salt from which you would like to prepare pure metallic sodium.

a) As the reduction potential of the aqueous Na⁺/Na couple is –1.662 V, this process is accomplished spontaneously in aqueous solution.

b) As the reduction potential of the aqueous Na⁺/Na couple is –1.662 V, this process is accomplished spontaneously from molten NaCl.

c) You prepare an aqueous solution of NaCl, apply a voltage of 1.662 V and collect metallic Na.

d) Na is produced by electrolysis of molten liquid NaCl at elevated temperature (NaCl mp. is 800^oC).

e) Na is produced by electrolysis of solid NaCl.

Difficulty: Medium

Learning Objective: Explain electrolytic reactions and cells.

Section Reference: 17.7 Electrolysis

Feedback: a) a large negative E^o value indicates that the species is not easily reduced, Na⁺ is not produced spontaneously from an aqueous solution of NaCl; b) as above, production of Na requires electrolysis; c) water is more easily reduced than is Na⁺, as a result you cannot produce Na from an aqueous solution of NaCl; d) correct answer; e) solid ionic NaCl is not conductive, electrolysis requires a conductive electrolyte

ESSAY QUESTIONS

29. Assign oxidation numbers to all the elements in HCO_­2­H.

Difficulty: Easy

Learning Objective: Use oxidation numbers to show what is being oxidized and what is being reduced in a redox reaction.

Section Reference: 17.1 Recognizing Redox Reactions

30. Assign oxidation numbers to all the elements in NO₂^-.

Difficulty: Easy

Learning Objective: Use oxidation numbers to show what is being oxidized and what is being reduced in a redox reaction.

Section Reference: 17.1 Recognizing Redox Reactions

31. Assign oxidation numbers to all the elements in titanium nitride, Ti₃N₄.

Difficulty: Easy

Learning Objective: Use oxidation numbers to show what is being oxidized and what is being reduced in a redox reaction.

Section Reference: 17.1 Recognizing Redox Reactions

32. Assign oxidation numbers to all the elements in sodium bicarbonate, NaHCO₃.

Difficulty: Easy

Learning Objective: Use oxidation numbers to show what is being oxidized and what is being reduced in a redox reaction.

Section Reference: 17.1 Recognizing Redox Reactions

33. Balance the following half reaction under neutral conditions:

HSO₃^- 🡪 SO₄^2-

Difficulty: Medium

Learning Objective: Balance redox reactions using the half-reaction method.

Section Reference: 17.2 Balancing Redox Reactions

34. Balance the following half reaction under acidic conditions:

I₂O₅ 🡪 I₂

Difficulty: Medium

Learning Objective: Balance redox reactions using the half-reaction method.

Section Reference: 17.2 Balancing Redox Reactions

35. Balance the following half reaction under basic conditions:

MnO₄^- 🡪 MnO_2(s)

Difficulty: Medium

Learning Objective: Balance redox reactions using the half-reaction method.

Section Reference: 17.2 Balancing Redox Reactions

36. Balance the following half reaction under basic conditions:

NO₃^- 🡪 NO₂^-

Difficulty: Medium

Learning Objective: Balance redox reactions using the half-reaction method.

Section Reference: 17.2 Balancing Redox Reactions

37. Balance the following half reaction under acidic conditions:

OCl^- 🡪 Cl^-

Difficulty: Medium

Learning Objective: Balance redox reactions using the half-reaction method.

Section Reference: 17.2 Balancing Redox Reactions

38. Use the half-reaction method to balance the following redox reaction:

Cl₂ 🡪 Cl- + ClO- (basic solution)

Difficulty: Medium

Learning Objective: Balance redox reactions using the half-reaction method.

Section Reference: 17.2 Balancing Redox Reactions

39. Use the half-reaction method to balance the following redox reaction:

I₂ + S₂O₃^2-🡪I - + S₄O₆^2- (basic solution)

Difficulty: Medium

Learning Objective: Balance redox reactions using the half-reaction method.

Section Reference: 17.2 Balancing Redox Reactions

40. If the coefficient of I^- is 1, determine the number of electrons transferred:

OCl^- + I^- 🡪 Cl^- + IO^-

Difficulty: Medium

Learning Objective: Balance redox reactions using the half-reaction method.

Section Reference: 17.2 Balancing Redox Reactions

41. Calculate the standard free energy change for the following redox reaction: (ΔG˚ = 77.11 (Ag⁺) and 65.49 kj/mol (Cu⁺²)

2 Ag⁺ (aq) + Cu (s) 🡪 2 Ag (s) + Cu²⁺ (aq)

Difficulty: Hard

Learning Objective: Describe galvanic cells.

Section Reference: 17.3 Galvanic Cells

42. Draw three molecular pictures illustrating direct electron transfer in the reaction of silver (I) ions with copper metal.

Difficulty: Medium

Learning Objective: Describe galvanic cells.

Section Reference: 17.3 Galvanic Cells

43. Draw a figure illustrating how a cell would be arranged for the redox reaction of copper with silver ion but using indirect electron transfer and a salt bridge with KNO₃ solution. Indicate the direction of electron flow in the wire and the movement of ions in the salt bridge.

Difficulty: Medium

Learning Objective: Describe galvanic cells.

Section Reference: 17.3 Galvanic Cells

44. Consider the redox process:

Cd(OH)₂ + Ni(OH)₂ + 2 OH^-  NiO₂ + Cd + H₂O

Write the equation for the spontaneous process and determine the free energy change for the spontaneous process.

Difficulty: Medium

Learning Objective: Describe galvanic cells.

Section Reference: 17.3 Galvanic Cells

45. For the galvanic cell shown in the diagram, identify the anode and mark which direction the cations are moving in the salt bridge.

Difficulty: Medium

Learning Objective: Describe galvanic cells.

Section Reference: 17.3 Galvanic Cells

46. Calculate the standard potential of voltaic cells that combine the following half reactions:

Pb to PbSO₄ and PbO₂ to PbSO₄ (acid solution)

(If needed, refer to Table 17-1.)

Difficulty: Medium

Learning Objective: Calculate standard cell potentials.

Section Reference: 17.4 Cell Potentials

47. Calculate the standard potential of the redox reaction:

2 Na + S 🡪 Na₂S (E° for S + 2e^– = S²- = –0.508 V

(If needed, refer to Table 17-1.)

Difficulty: Easy

Learning Objective: Calculate standard cell potentials.

Section Reference: 17.4 Cell Potentials

48. Calculate the standard potential of the aluminium air battery in which the active materials Al(s) andO₂, and the electrolyte is aqueous KOH.

(If needed, refer to Table 17-1.)

Difficulty: Hard

Learning Objective: Calculate standard cell potentials.

Section Reference: 17.4 Cell Potentials

Feedback: Student must recognize galvanic cell with Al³⁺/Al couple, E^o=–1.662 and O₂/OH^- couple with E^o = 0.401, give E­^o_cell = 2.063 = 0.401–(–1.662)

49. Balance the reaction and calculate the standard potential for:

Zr + H₂O 🡪 ZrO₂ + H₂

Given: ZrO₂ + 4e^- + 4 H₃O⁺ 🡪 Zr + 6 H₂O (E°=–1.43 V)

(If needed, refer to Table 17-1.)

Difficulty: Medium

Learning Objective: Calculate standard cell potentials.

Section Reference: 17.4 Cell Potentials

50. An electrochemical cell is constructed that contains Cr³⁺(aq) and Cr metal as the electrode in one compartment and Cu²⁺(aq) and copper metal in the other compartment. Calculate the expected standard potential upon appropriately connecting the cell and describe the direction of electron and cation flow.

(If needed, refer to Table 17-1.)

Difficulty: Medium

Learning Objective: Calculate standard cell potentials.

Section Reference: 17.4 Cell Potentials

51. Which of the species listed is the strongest oxidizing agent?

Half-cell reaction E ^o (v)

Cu²⁺ (aq) + 2e^- Cu (s) +0.342

Fe²⁺ (aq) + 2e^- Fe (s) –0.45

Cr³⁺ (aq) + 2e^- Cr (s) –0.744

Al³⁺ (aq) + 3e^- Al (s) –1.66

Pt²⁺ (aq) + 2e^- Pt (s) +1.20

(If needed, refer to Table 17-1.)

Difficulty: Easy

Learning Objective: Calculate standard cell potentials.

Section Reference: 17.4 Cell Potentials

52. Which of the species listed is the strongest reducing agent?

Half-cell reaction E ^o (v)

Cu²⁺ (aq) + 2e^- Cu (s) +0.342

Fe²⁺ (aq) + 2e^- Fe (s) –0.45

Cr³⁺ (aq) + 2e^- Cr (s) –0.744

Al³⁺ (aq) + 3e^- Al (s) –1.66

Pt²⁺ ₍aq) + 2e^- Pt (s) +1.20

(If needed, refer to Table 17-1.)

Difficulty: Easy

Learning Objective: Calculate standard cell potentials.

Section Reference: 17.4 Cell Potentials

53. Calculate the standard free energy change for the redox reaction between silver ion and copper to give copper (II) and silver metal.

(If needed, refer to Table 17-1.)

Difficulty: Hard

Learning Objective: Relate cell potential to the reaction conditions.

Section Reference: 17.5 Free Energy and Electrochemistry

54. Calculate the standard free energy changes for the following redox reaction:

2 Ag⁺_(aq) + Sn²⁺_(aq) 🡪 2 Ag_(s) + Sn⁴⁺_(aq) [E˚(Sn^{4+, 2+}) = 0.151 V]

(If needed, refer to Table 17-1.)

Difficulty: Medium

Learning Objective: Relate cell potential to the reaction conditions.

Section Reference: 17.5 Free Energy and Electrochemistry

55. Calculate the equilibrium constant for the following redox reaction:

Fe3+ (aq) + Cu+ (aq) 🡪 Fe2+ (aq) + Cu2+ (aq)

[E˚(Fe^{3+, 2+}) = 0.771 V] [E˚(Cu^2+,1+) = 0.153 V]

(If needed, refer to Table 17-1.)

Difficulty: Hard

Learning Objective: Relate cell potential to the reaction conditions.

Section Reference: 17.5 Free Energy and Electrochemistry

56. Calculate the equilibrium constants for the following redox reaction:

2 Cu²+ (aq) + Sn2+ (aq) 🡪 2 Cu+ (aq) + Sn⁴+ (aq)

[E˚(Sn^{4+, 2+}) = 0.151 V] [E˚(Cu^{2+, 1+}) = 0.153 V]

(If needed, refer to Table 17-1.)

Difficulty: Hard

Learning Objective: Relate cell potential to the reaction conditions.

Section Reference: 17.5 Free Energy and Electrochemistry

57. Consider the Daniell cell for which the cell reaction and standard potential are:

Zn(s) + Cu²⁺ (aq) 🡪 Zn²⁺ (aq) + Cu (s) E° = 1.10 V

If the cell is initially at standard conditions ([Cu²⁺] = [Zn²⁺] = 1.00 M) and assuming it contains 1 L of electrolyte, determine the mass of Zn(s) lost when the cell potential falls to 1.06 V?

(If needed, refer to Table 17-1.)

Difficulty: Hard

Learning Objective: Relate cell potential to the reaction conditions.

Section Reference: 17.5 Free Energy and Electrochemistry

58. Consider an electrochemical cell of the type shown in the figure where the redox half-reaction in both compartments has the identical standard potentials:

Use the Nernst equation to calculate the potential developed by this cell.

(If needed, refer to Table 17-1.)

Difficulty: Hard

Learning Objective: Relate cell potential to the reaction conditions.

Section Reference: 17.5 Free Energy and Electrochemistry

59. Aluminium is used in a battery in which the following reaction occurs:

4 Al (s) + 3 O₂ (g) + 4 OH- (aq) + 6 H₂O 🡪4 Al(OH)

If the battery must supply a current of 78 A for 4.0 hours, what mass of Al (in g) must be contained in the battery?

(If needed, refer to Table 17-1.)

Difficulty: Hard

Learning Objective: Relate cell potential to the reaction conditions.

Section Reference: 17.5 Free Energy and Electrochemistry

60. An electrochemical cell is made by immersing a piece of Cd metal into a solution of 0.100 M CdSO₄ and a Zn electrode into a solution of 1.00 M ZnSO₄ and placing a salt bridge to allow ion flow between the two solutions. a) What voltage will be produced by the cell and b) what metal is the anode? (Cd²⁺ + 2e^-🡪 Cd; E° = –0.402 V)

(If needed, refer to Table 17-1.)

Difficulty: Medium

Learning Objective: Relate cell potential to the reaction conditions.

Section Reference: 17.5 Free Energy and Electrochemistry

61. The lead–acid battery used in automobiles utilizes the following redox reaction:

PbO₂(s) + Pb (s) + 2 HSO₄^- (aq) + 2 H₃O⁺ (aq) 🡪 2 PbSO₄(s) + 4 H₂O (l) E°= 2.04 V

What mass of H₂ being oxidized by O₂ under standard acid conditions would be required to give the same amount of electrons as one mole of lead oxide?

(If needed, refer to Table 17-1.)

Difficulty: Hard

Learning Objective: Relate cell potential to the reaction conditions.

Section Reference: 17.5 Free Energy and Electrochemistry

62. Consider an automobile which is powered by a perfectly efficient fuel cell that consumes hydrogen and oxygen in the following redox reaction:

2H₂(g) + O₂(g) 🡪 2 H₂O (l)

If the electric system requires a current of 500 amperes, how many g of H₂ are consumed per hour?

(If needed, refer to Table 17-1.)

Difficulty: Hard

Learning Objective: Relate cell potential to the reaction conditions.

Section Reference: 17.5 Free Energy and Electrochemistry

63. For the working galvanic cell shown at standard conditions, determine the balanced reaction and direction of electron flow through the wire.

(If needed, refer to Table 17-1.)

Difficulty: Medium

Learning Objective: Relate cell potential to the reaction conditions.

Section Reference: 17.5 Free Energy and Electrochemistry

64. The same charge of 1.07 x 10⁴ C is passed through three solutions: one each of Au³⁺, Cu⁺ and Pb²⁺ with strips of the metals as cathodes. In which cell will the greatest mass of metal be reduced and what is the mass of that metal?

Difficulty: Medium

Learning Objective: Explain electrolytic reactions and cells.

Section Reference: 17.7 Electrolysis

65. For a brine electrolysis cell (see redox reaction below) operating at 60,000 amps, how many kg of NaOH and Cl₂ would be produced in 24.0 hours?

2 NaCl (aq) + 2 H₂O 🡪 2 NaOH (aq) + Cl₂(g) + H₂ (g)

Difficulty: Medium

Learning Objective: Explain electrolytic reactions and cells.

Section Reference: 17.7 Electrolysis

66. An electrolytic cell driving the following redox reaction has a current of 4.02 amps passed through it for 2.32 hours. How much Ag will be dissolved and how much Cu will be deposited?

2 Ag (s) + Cu⁺² (aq) 🡪2 Ag⁺ (aq) + Cu (s)

Difficulty: Medium

Learning Objective: Explain electrolytic reactions and cells.

Section Reference: 17.7 Electrolysis

67. Determine what causes the following electrolytic cell (which includes 50 g of metallic Ag and 1 L of 0.15 M Cu(NO₃)₂ ) to cease operation and determine how long the cell can sustain a current of 5 amps.

2 Ag (s) + Cu⁺² (aq) 🡪2 Ag⁺ (aq) + Cu (s)

Difficulty: Hard

Learning Objective: Explain electrolytic reactions and cells.

Section Reference: 17.7 Electrolysis

68. At an engine block rebuilding factory you are in charge of replating Mn on the interiors of engine blocks. Based on the surface area and thickness needed, you determine that you need 35g of Mn to plate out by performing electrolysis on the engine block. Your plating solution is 3 M Mn(NO₃)₂. How long do you need to perform electrolysis if your machine performs at 220 Amps?

Difficulty: Hard

Learning Objective: Explain electrolytic reactions and cells.

Section Reference: 17.7 Electrolysis

69. You determine that for proper protection of an engine part you need to put a coating of 3.0g of Cr_(s) on your part. How long do you need to perform electrolysis on your engine part if your current is 30.0 Amps, and your Cr is in the form of Cr(NO₃)_3(aq)?

Difficulty: Medium

Learning Objective: Explain electrolytic reactions and cells.

Section Reference: 17.7 Electrolysis

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