Deck 17: Electron Transfer Reactions

Question 1

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

Answer 1

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

Question 2

Balance redox reactions using the half-reaction method.

Answer 2

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

Question 3

Describe galvanic cells.

Answer 3

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

Question 4

Calculate standard cell potentials.

Question 5

Relate cell potential to the reaction conditions.

Question 6

Explain the chemistry of everyday redox reactions.

Question 7

Explain electrolytic reactions and cells.

Question 8

Which of the following is NOT a redox reaction?

A) 4 NH₃ + 5 O₂ $\rarr$ 4 NO + 6 H₂O
B) 2 CO + O₂ $\rarr$ 2 CO₂
C) S + 2 F₂ $\rarr$ SF₄
D) AgNO₃ + KI $\rarr$ AgI(s) + KNO₃
E) Cl₂ + 2H₂O (\rarr\) 2Cl^- + 2OCl^- + 4H⁺

Question 9

Which of the following are redox reactions?
I. 2 N₂ + 3 H₂ $\rarr$ 2 NH₃
II. 4 Al + 3 O₂ $\rarr$ 2 Al₂O₃
III. 2 NO₂ $\rarr$ N₂O₄
IV. FeCl₃ (aq) + 6 NH₃(aq) (\rarr\)Fe(NH₃)₆³⁺ (aq) + 3 Cl^- (aq)
V. Cu²⁺ + Zn $\rarr$ 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

Question 10

Which elements are changing oxidation states in the following reaction?
Zn(s) + 2 MnO₂(s) + H₂O(l) $\rarr$ Zn(OH)₂(s) + Mn₂O₃(s)

A) Zn, O
B) Mn, Zn
C) H, Mn
D) O, Mn
E) H, O

Question 11

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

Question 12

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

Question 13

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

Question 14

Determine the coefficient for Sn⁺² in the following balanced redox reaction.MnO₄^- + Sn²⁺ $\rarr$ Sn⁴⁺ + Mn²⁺ (acidic solution)

A) 2
B) 4
C) 5
D) 6
E) 10

Question 15

Consider the redox reaction of permanganate and sulphur:
MnO₄^- + S $\rarr$ 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

Question 16

Consider the redox reaction of nitric acid and copper:
Cu + HNO₃ $\rarr$ 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

Question 17

Consider the redox reaction of triiodide and oxygen:I₃^- + O₂ $\rarr$ 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

Question 18

The following reaction occurs in a galvanic cell:
NiO₂ + Cd + H₂O $\rarr$ Cd(OH)₂ + Ni(OH)₂ + 2 OH^-Which redox process in this battery occurs at a passive electrode?

A) Cd $\rarr$ Cd(OH)₂
B) NiO₂ $\rarr$ Ni(OH)₂
C) O₂ $\rarr$ 4 OH^-
D) H₂O $\rarr$ OH^-
E) neither electrode is passive

Question 19

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

Question 20

For the reaction given, which half reaction occurs at the cathode?
NiO₂ + Cd + H₂O $\rarr$ Cd(OH)₂ + Ni(OH)₂ + 2 OH^-

A) Cd $\rarr$ Cd(OH)₂
B) NiO₂ $\rarr$ Ni(OH)₂
C) Cd(OH)₂ $\rarr$ Cd
D) H₂O $\rarr$ OH^-
E) There is no cathode as the cathodic reaction always occurs at the passive electrode.

Question 21

For the reaction given below, which half reaction occurs at the cathode?Pb(s) + PbO₂(s) + 2HSO₄^-(aq) + 2H₃O⁺(aq) $\rarr$ 2PbSO₄(aq) +4H₂O(l)

A) Pb $\rarr$ PbSO₄
B) PbO₂ $\rarr$ PbSO₄
C) HSO₄^- $\rarr$ PbSO₄
D) H₃O⁺ $\rarr$ H₂O
E) PbSO₄ $\rarr$ PbO₂

Question 22

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) $\rarr$ 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

Question 23

For the reaction given below, which half reaction occurs at the anode?
2 H₂(g) + O₂(g) $\rarr$ 2 H₂O(l)

A) H₂ $\rarr$ H₂O
B) O₂ $\rarr$ H₂O
C) H₂O $\rarr$ H₂
D) H₂O $\rarr$ O₂
E) H₂(g) $\rarr$ 2H⁺

Question 24

Consult a table of reduction potentials (Table 17-1 in the text) 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

Question 25

Platinum metal is quite resistant to oxidation as may be deduced by its reduction potential:Pt²⁺ + 2e $\rarr$ $\quad$ Pt E° $\thickapprox$ 1.2 VExamine a table of reduction potentials (Table 17-1 in the text) 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

Question 26

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 in the text )

A) 0.259 V, Fe(s) + 3AgCl(s) $\rarr$ Fe³⁺(aq) + 3Ag(s) + 3Cl^-(aq)
B) 0.669 V, Fe(s) + 2AgCl(s) $\rarr$ Fe²⁺(aq) + 2Ag(s) + 2Cl^-(aq)
C) -0.669 V, Fe(s) + 2AgCl(s) $\rarr$ Fe²⁺(aq) + 3Ag(s) + 3Cl^-(aq)
D) -0.225 V, Fe(s) + 2AgCl(s) $\rarr$ Fe²⁺(aq) + 2Ag(s) + 2Cl^-(aq)
E) 0.669 V, Fe(s) + AgCl(s) $\rarr$ Fe²⁺(aq) + Ag(s) + Cl^-(aq)

Question 27

Consider the Daniell cell where the cell reaction and standard potential are:Zn(s) + Cu²⁺ (aq) $\rarr$ Zn²⁺ (aq) + Cu (s) $\quad$ E° = 1.10 VIf 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: $\Delta$ G° = -nFE°, E° , E = E° - , moles e^- = (If needed, refer to Table 17-1in the text)

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

Question 28

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 $\Delta$ G = -2 kJ?If needed, use the following equation: $\Delta$ G° = -nFE°, E° , E = E?
- , moles e^- = (If needed, refer to Table 17-1 in the text )

A) Ag(s) $\mid$ Ag⁺( aq, 1.00 M) $\parallel$ Ag⁺(aq, 2.25 M) $\mid$ Ag(s)
B) Ag(s) $\mid$ Ag⁺( aq, 1.00 M) $\parallel$ Ag⁺(aq, 1.00 M) $\mid$ Ag(s)
C) Ag(s) $\mid$ Ag⁺( aq, 2.25 M) $\parallel$ Ag⁺(aq, 1.00 M) $\mid$ Ag(s)
D) Ag(s) $\mid$ Ag⁺( aq, 0.445 M) $\parallel$ Ag⁺(aq, 1.00 M) $\mid$ Ag(s)
E) Ag(s) $\mid$ Ag⁺( aq, 1.00 M), Ag⁺(aq, 2.25 M) $\mid$ Ag(s)

Question 29

For the working galvanic cell shown at standard conditions, how would you increase the cell potential?(If needed, refer to Table 17-1 in the text )

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.

Question 30

For the following galvanic cell what will be its potential when the reaction reaches equilibrium?(If needed, refer to Table 17-1in the text )

A) 0.0 V
B) 0.458 V
C) 1.142 V
D) 0.272 V
E) 1.26 V

Question 31

Ships, storage tanks, and other large metal items may be protected from corrosion by(If needed, refer to Table 17-1in the text )

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.

Question 32

What are the possible oxidation states of corroded iron?(If needed, refer to Table 17-1 in the text)

A) 2
B) 0
C) 3
D) 2 and 3
E) 0, 2 and 3

Question 33

How is aluminium protected from oxidation?(If needed, refer to Table 17-1in the text)

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.

Question 34

Which of the following combinations would provide the largest potential for a battery?(If needed, refer to Table 17-1in the text)

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^-

Question 35

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.

Question 36

Assign oxidation numbers to all the elements in HCO₂H.

Question 37

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

Question 38

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

Question 39

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

Question 40

Balance the following half reaction under neutral conditions:HSO₃^- $\rightarrow$ SO₄^2-

Question 41

Balance the following half reaction under acidic conditions:I₂O₅ $\rightarrow$ I₂

Question 42

Balance the following half reaction under basic conditions:
MnO₄^- $\rightarrow$ MnO_2(s)

Question 43

Balance the following half reaction under basic conditions:NO₃^- $\rightarrow$ NO₂^-

Question 44

Balance the following half reaction under acidic conditions:OCl^- $\rightarrow$ Cl^-

Question 45

Use the half-reaction method to balance the following redox reaction:
Cl₂ $\rightarrow$ Cl^- + ClO^- (basic solution)

Question 46

Use the half-reaction method to balance the following redox reaction:
I₂ + S₂O₃^2- $\rightarrow$ I ^- + S₄O₆^2- (basic solution)

Question 47

If the coefficient of I^- is 1, determine the number of electrons transferred:
OCl^- + I^- $\rightarrow$ Cl^- + IO^-

Question 48

Calculate the standard free energy change for the following redox reaction: ( $\Delta$ G $\degree$ = 77.11 (Ag⁺) and 65.49 kj/mol
(Cu⁺²)2 Ag⁺ (aq) + Cu (s) $\rightarrow$ 2 Ag (s) + Cu²⁺ (aq)

Question 49

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

Question 50

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.

Question 51

Consider the redox process:
Cd(OH)₂ + Ni(OH)₂ + 2 OH^- $\rightarrow$ NiO₂ + Cd + H₂OWrite the equation for the spontaneous process and determine the free energy change for the spontaneous process.

Question 52

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

Question 53

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.in the text )

Question 54

Calculate the standard potential of the redox reaction:
2 Na + S $\rightarrow$ Na₂S (E° for S + 2e^- = S^2- = -0.508 V(If needed, refer to Table 17-1. in the text)

Question 55

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 in the text)

Question 56

Balance the reaction and calculate the standard potential for:
Zr + H₂O $\rightarrow$ ZrO₂ + H₂Given: ZrO₂ + 4e^- + 4 H₃O⁺ $\rightarrow$ Zr + 6 H₂O (E°=-1.43 V)(If needed, refer to Table 17-1 in the text)

Question 57

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 in the text)

Question 58

Which of the species listed is the strongest oxidizing agent?‪ (If needed, refer to Table 17-1 in the text)

Question 59

Which of the species listed is the strongest reducing agent? (If needed, refer to Table 17-1 in the text )

Question 60

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 in the text)

Question 61

Calculate the standard free energy changes for the following redox reaction:
2 Ag⁺_(aq) + Sn²⁺_(aq) $\rightarrow$ 2 Ag_(s) + Sn⁴⁺_(aq) [E°(Sn^{4+, 2+}) = 0.151 V](If needed, refer to Table 17-1 in the text)

Question 62

Calculate the equilibrium constant for the following redox reaction:
Fe³⁺ (aq) + Cu⁺ (aq) $\rightarrow$ 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 in the text)

Question 63

Calculate the equilibrium constants for the following redox reaction:
2 Cu²⁺ (aq) + Sn²⁺ (aq) $\rightarrow$ 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 in the text)

Question 64

Consider the Daniell cell for which the cell reaction and standard potential are:Zn(s) + Cu²⁺ (aq) $\rightarrow$ Zn²⁺ (aq) + Cu (s) E° = 1.10 VIf 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 17-1in the text)

Question 65

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 in the text)

Question 66

Aluminium is used in a battery in which the following reaction occurs:
4 Al (s) + 3 O₂ (g) + 4 OH^- (aq) + 6 H₂O $\rightarrow$ 4 Al(OH) (-,4) (aq)If the battery must supply a current of 78 A for 4.0 hours, what mass of Al (ing) must be contained in the battery?(If needed, refer to Table 17-1 in the text)

Question 67

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^- $\rightarrow$ Cd; E° = -0.402 V)(If needed, refer to Table 17-1 in the text)

Question 68

The lead-acid battery used in automobiles utilizes the following redox reaction:
PbO₂(s) + Pb (s) + 2 HSO₄^- (aq) + 2 H₃O⁺ (aq) $\rightarrow$ 2 PbSO₄(s) + 4 H₂O (l) E°= 2.04 VWhat 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 in the text)

Question 69

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) $\rightarrow$ 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 in the text )

Question 70

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 in the text )

Question 71

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?

Question 72

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 $\rightarrow$ 2 NaOH (aq) + Cl₂(g) + H₂ (g)

Question 73

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) $\rightarrow$ 2 Ag⁺ (aq) + Cu (s)

Question 74

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) $\rightarrow$ 2 Ag⁺ (aq) + Cu (s)

Question 75

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?

Question 76

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)?