Homework answers / question archive / Chapter 19: Electrochemistry 1)Complete and balance the following redox equation

Chapter 19: Electrochemistry 1)Complete and balance the following redox equation

Chemistry

Share With

---

Chapter 19: Electrochemistry

1)Complete and balance the following redox equation.  When properly balanced using the smallest whole-number coefficients, the coefficient of S is

            H₂S + HNO₃ → S + NO  (acidic solution)

            A)  1    B)  2    C)  3    D)  5    E)  6

                                                      

 

2. Complete and balance the following redox equation using the smallest whole-number coefficients.  What is the coefficient of Sn in the balanced equation?

            Sn + HNO₃ → SnO₂ + NO₂ + H₂O  (acidic solution)

            A)  1    B)  2    C)  3    D)  4    E)  5

                                            Ans:  A    

 

3. Complete and balance the following redox equation.  What is the coefficient of H₂O when the equation is balanced using the set of smallest whole-number coefficients?

  MnO₄^– + SO₃^2– → Mn²⁺ + SO₄^2–  (acidic solution)   A)  3    B)  4    C)  5    D)  8    E)  None of these.

                                            Ans:  A    

 

4. Complete and balance the following redox equation.  What is the coefficient of H₂S when the equation is balanced using the set of smallest whole-number coefficients?   H₂S + MnO₄^– → Mn²⁺ + SO₄^2–  (acidic solution)   A)  1    B)  2    C)  4    D)  5    E)  None of these.

                                                       

 

5. Complete and balance the following redox equation.  What is the coefficient of H₂O when the equation is balanced with the set of smallest whole-number coefficients?   H₂O + MnO₄^– + I^– → MnO₂ + IO₃^–  (basic solution)   A)  1    B)  2    C)  4    D)  10    E)  None of these.

                                            Ans:  A    

 

6. Complete and balance the following redox equation.  What is the coefficient of OH^– when the equation is balanced using the set of smallest whole-number coefficients?

             MnO₄^– + I^– → MnO₂ + IO₃^–  (basic solution)

            A)  1    B)  2    C)  4    D)  10    E)  None of these.

                                                      

 

7. Complete and balance the following redox equation using the set of smallest wholenumber coefficients.  Now sum the coefficients of all species in the balanced equation. 

(Remember the coefficients that are equal to one.)  The sum of the coefficients is  BrO₃^–(aq) + Sb³⁺(aq) → Br^–(aq) + Sb⁵⁺(aq)  (acidic solution)

 A)  4    B)  12    C)  13    D)  17    E)  None of these.

                                                       

 

8. Complete and balance the following redox equation that occurs in acidic solution using the set of smallest whole-number coefficients.  What is the sum of all the coefficients in the equation?

             PbO₂(s) + Cl^– → Pb²⁺ + Cl₂(g)  (acidic solution)

            A)  2    B)  4    C)  5    D)  9    E)  11

                                                      

 

9. Complete and balance the following redox equation using the set of smallest wholenumbers coefficients.  What is the sum of the coefficients?

            HI + HNO₃ → I₂ + NO  (acidic solution)

            A)  5    B)  7    C)  14    D)  17    E)  None of these.

                                                       

 

10. Complete and balance the following redox equation.  The sum of the smallest wholenumber coefficients is

             MnO₄^– + H⁺ + Br^– → Mn²⁺ + Br₂ + H₂O  (acidic solution)

            A)  6    B)  17    C)  21    D)  29    E)  43

                                                      

 

11. Complete and balance the following redox equation.  The sum of the smallest wholenumber coefficients is

             Bi(OH)₃ + SnO₂^2– → Bi + SnO₃^2–  (basic solution)

            A)  32    B)  25    C)  16    D)  13    E)  None of these.

                                                       

 

12. Complete and balance the following redox equation.  The sum of the smallest wholenumber coefficients is

             Br₂ → BrO₃^– + Br^–  (basic solution)

            A)  9    B)  12    C)  18    D)  21    E)  None of the above.

                                                      

 

13. Given the following notation for an electrochemical cell  Pt(s) | H₂(g) | H⁺(aq) || Ag⁺(aq) | Ag(s)  what is the balanced overall (net) cell reaction?
    1. 1. 2H⁺(aq) + 2Ag⁺(aq) → H₂(g) + 2Ag(s)
       2. H₂(g) + 2Ag(s) → H⁺(aq) + 2Ag⁺(aq)
       3. 2H⁺(aq) + 2Ag(s) → H₂(g) + 2Ag⁺(aq)
       4. H₂(g) + Ag⁺(aq) → H⁺(aq) + Ag(s)
       5. H₂(g) + 2Ag⁺(aq) → 2H⁺(aq) + 2Ag(s)

                                                        

 

14. Consider an electrochemical cell constructed from the following half cells, linked by an external circuit and by a KCl salt bridge.
    • • an Al(s) electrode in 1.0 M Al(NO₃)₃ solution
      • a Pb(s) electrode in 1.0 M Pb(NO₃)₂ solution

             The balanced overall (net) cell reaction is

1. 1. 1. Pb(s) + Al³⁺(aq) → Pb²⁺(aq) + Al(s.)
      2. 3Pb(s) + 2Al³⁺(aq) → 3Pb²⁺(aq) + 2Al(s).
      3. 3Pb²⁺(aq) + 2Al(s) → 3Pb(s) + 2Al³⁺(aq).
      4. Pb²⁺(aq) + Al(s) → Pb(s) + Al³⁺(aq).

                                                        

 

15. Consider an electrochemical cell constructed from the following half cells, linked by a KCl salt bridge.

  • an Al(s) electrode in 0.5 M Al₂(SO₄)₃ solution   • a Pb(s) electrode in 1.0 M Pb(NO₃)₂ solution   Which electrode is the anode?

            A)  Al    B)  Pb    C)  Neither

                                                   

 

16. Consider an electrochemical cell constructed from the following half cells, linked by a KCl salt bridge.
    • • a Fe electrode in 1.0 M FeCl₂ solution
      • a Sn electrode in 1.0 M Sn(NO₃)₂ solution

             When the cell is running spontaneously, which choice includes only true statements and no false ones?

1. 1. 1. The tin electrode loses mass and the tin electrode is the cathode.
      2. The tin electrode gains mass and the tin electrode is the cathode.
      3. The iron electrode gains mass and the iron electrode is the anode.
      4. The iron electrode loses mass and the iron electrode is the cathode.
      5. The iron electrode gains mass and the iron electrode is the cathode.

                                                        

 

17. Consider an electrochemical cell constructed from the following half cells, linked by a KCl salt bridge.
    • • a Fe electrode in 1.0 M FeCl₂ solution
      • a Ni electrode in 1.0 M Ni(NO₃)₂ solution

 When the cell is running spontaneously, which choice includes only true statements and no false ones?

1. 1. 1. The nickel electrode loses mass and the nickel electrode is the cathode.
      2. The nickel electrode gains mass and the nickel electrode is the cathode.
      3. The iron electrode gains mass and the iron electrode is the anode.
      4. The iron electrode loses mass and the iron electrode is the cathode.

                                                        

 

18. A certain electrochemical cell has for its cell reaction:

            Zn + HgO → ZnO + Hg

            Which is the half-reaction occurring at the anode?

1. 1. 1. HgO + 2e^– → Hg + O^2–       C)        Zn → Zn²⁺ + 2e^–
      2. Zn²⁺ + 2e^– → Zn      D)        ZnO + 2e^– → Zn

                                                        

 

19. Calculate the value of E°_cell for the following reaction:

             2Au(s) + 3Ca²⁺(aq) → 2Au³⁺(aq) + 3Ca(s)

            A)  –4.37 V    B)  –1.37 V    C)  –11.6 V    D)  1.37 V    E)  4.37 V

                                                   

 

20. Calculate E°_cell for a silver-aluminum cell in which the cell reaction is

             Al(s) + 3Ag⁺(aq) → Al³⁺(aq) + 3Ag(s)

            A)  –2.46 V    B)  0.86 V    C)  –0.86 V    D)  2.46 V    E)  none of these

                                                         

 

21. Calculate E°_cell for the following reaction:

             2Fe²⁺(aq) + Cd²⁺(aq) → 2Fe³⁺(aq) + Cd(s)

            A)  –0.37 V    B)  0.37 V    C)  –1.17 V    D)  1.17 V    E)  None of these.

                                                        

 

22. For the reaction, 2Cr²⁺ + Cl₂(g) → 2Cr³⁺ + 2Cl^–, E°_cell  is 1.78 V. Calculate E°_cell for the related reaction Cr³⁺ + Cl^– → Cr²⁺ + ¹/₂Cl₂(g).

             A)  1.78 V    B)  0.89 V    C)  –1.78 V    D)  –0.89 V    E)  None of these.                    

 

23. For the reaction Ni²⁺(aq) + 2Fe²⁺(aq) → Ni(s) + 2Fe³⁺(aq), the standard cell potential

E°_cell is

            A)  +2.81 V.    B)  +1.02 V.    C)  +0.52 V.    D)  –1.02 V.    E)  –2.81 V.

                                                         

 

24. Calculate the standard cell emf for the following cell:

 Mg | Mg²⁺ || NO₃^– (acid soln)| NO(g) | Pt

 A)  3.33 V    B)  1.41 V    C)  –1.41 V    D)  8.46 V    E)  –8.46 V 

25. Consider an electrochemical cell based on the following cell diagram:

             Pt | Pu³⁺(aq), Pu⁴⁺(aq) || Cl₂(g), Cl^– (aq) | Pt

  Given that the standard cell emf is 0.35 V and that the standard reduction potential of chlorine is 1.36 V, what is the standard reduction potential E°(Pu⁴⁺/Pu³⁺)?   A)  2.37 V    B)  1.01 V    C)  –1.71 V    D)  –1.01 V    E)  1.71 V

                                                        

 

26. Consider the following electrochemical cell:

             U | U³⁺(aq) || Cl^–(aq),Cl₂(g) | Pt

            If the standard cell emf is 3.16 V, what is the standard reduction potential for uranium?

             A)  –3.16 V    B)  +3.16 V    C)  –1.80 V    D)  +1.80 V    E)  +1.36 V              

 

27. The overall reaction 2Co³⁺(aq) + 2Cl^–(aq) → 2Co²⁺(aq) + Cl₂(g) has the standard cell voltage E°_cell= 0.46 V. Given E° = 1.36 V for the reaction Cl₂(g) + 2e^– → 2Cl^–(aq), calculate the standard reduction potential for the following the half reaction at 25°C:

             Co³⁺ + e^– → Co²⁺

            A)  1.82 V    B)  –0.90 V    C)  0.90 V    D)  –1.82 V    E)  –1.36 V

                                                   

 

28. An electrochemical cell based on the following reaction has a standard cell voltage

(E°_cell) of 0.48 V:

             Sn(s) + Cu²⁺(aq) → Sn²⁺(aq) + Cu(s)

            What is the standard reduction potential of tin(II)?  (E°(Cu²⁺/Cu) = 0.34 V)

            A)  –0.14 V    B)  0.14 V    C)  –0.82 V    D)  0.82 V    E)  none of these

                                                   

 

29. Consider a voltaic cell based on the following cell reaction:

             Ni(s) + At₂(s) → Ni²⁺(aq) + 2At^– (aq)

             Given that the standard cell emf is 0.55 V, what is the standard reduction potential for astatine? [E°(Ni²⁺/Ni) = –0.25 V]

            A)  0.80 V    B)  0.30 V    C)  –0.30 V    D)  –0.80 V    E)  0.43 V

                                                        

 

30. According to the following cell diagram, which chemical species undergoes reduction?

            Sn | Sn²⁺ || NO₃^– (acid soln), NO(g) | Pt

            A)  Sn    B)  Sn²⁺    C)  NO₃^–    D)  NO    E)  Pt

                                                        

 

31. In the following half equation, which is the oxidizing agent?

 NO₃^–(aq) + 4H⁺(aq) + 3e^– → NO(g) + 2H₂O

 A)  NO₃^–    B)  H⁺    C)  e^–    D)  NO    E)  H₂O

 Ans:  A    

32. Which statement is true for a spontaneous redox reaction carried out at standard-state conditions?
    1. 1. E°_red is always negative.      C)        E°_ox is always positive.
       2. E°_cell is always positive.       D)        E°_red is always positive.

                       Category:  Easy     Section:  19.3

 

33. Which one of the following reactions will occur spontaneously at standard-state conditions and 25°C?
    1. 1. Mg²⁺ + Ca → Mg + Ca²⁺     C)        2Al³⁺ + 3Fe → 2Al + 3Fe²⁺
       2. Au + 3K⁺ → Au³⁺ + 3K      D)        Cu + 2H⁺ → Cu²⁺ + H₂

                                                   

 

34. Consider the following standard reduction potentials in acid solution:

             

                                                                                E°(V)

                  Al³⁺ + 3e^– → Al(s)                                 –1.66

                  AgBr(s) + e^– → Ag(s) + Br^–                  +0.07

                  Sn⁴⁺ + 2e^– → Sn²⁺                                  +0.14

                  Fe³⁺ + e^– → Fe²⁺                                    +0.77

 

  The strongest reducing agent among those shown above is   A)  Fe³⁺.    B)  Fe²⁺.    C)  Br^–.    D)  Al³⁺.    E)  Al.

                                                        

 

35. Consider the following standard reduction potentials in acid solution:

             

                                                                                E°(V)

                  Al³⁺ + 3e^– → Al(s)                                 –1.66

                  AgBr(s) + e^– → Ag(s) + Br^–                  +0.07

                  Sn⁴⁺ + 2e^– → Sn²⁺                                  +0.14

                  Fe³⁺ + e^– → Fe²⁺                                    +0.77

 

  The strongest oxidizing agent among those shown above is   A)  Fe³⁺.    B)  Fe²⁺.    C)  Br^–.    D)  Al³⁺.    E)  Al.

                                                   

 

36. Consider the following standard reduction potentials in acid solution:

 

	E°(V)
Al3+ + 3e– → Al(s)	–1.66
Sn4+ + 2e– → Sn2+	+0.14
I2(s) + 2e– → 2I– (aq)	+0.53

 

            Which is the weakest oxidizing agent in this list?

             A)  Al³⁺(aq)    B)  Al(s)    C)  I^–(aq)    D)  I₂(s)    E)  Sn⁴⁺(aq)

                                                   

 

37. Consider the following standard reduction potentials in acid solution:

                  Cr³⁺ + 3e^– → Cr                                           E° = –0.74 V

                  Co²⁺ + 2e^– →  Co                                         E° = –0.28 V

                  MnO₄^–  + 8H⁺ + 5e^– → Mn²⁺ + 4H₂O          E° = +1.51 V

 

             The weakest reducing agent listed above is

            A)  Cr³⁺.    B)  Cr.    C)  Mn²⁺.    D)  Co.    E)  MnO₄^–.

                                                        

 

38. Consider the following standard reduction potentials in acid solution:

                  Cr³⁺ + 3e^– → Cr                                           E° = –0.74 V

                  Co²⁺ + 2e^– →  Co                                         E° = –0.28 V

                  MnO₄^–  + 8H⁺ + 5e^– → Mn²⁺ + 4H₂O          E° = +1.51 V

 

             The strongest reducing agent listed above is

            A)  Cr³⁺.    B)  Cr.    C)  Mn²⁺.    D)  Co.    E)  MnO₄^–.

                                                        

 

39. Consider the following standard reduction potentials in acid solution:

                  Cr³⁺ + 3e^– → Cr                                           E° = –0.74 V

                  Co²⁺ + 2e^– →  Co                                         E° = –0.28 V

                  MnO₄^–  + 8H⁺ + 5e^– → Mn²⁺ + 4H₂O          E° = +1.51 V

 

             The strongest oxidizing agent listed above is

             A)  Cr³⁺.    B)  Cr.    C)  Mn²⁺.    D)  Co²⁺.    E)  MnO₄^–.

                                                        

 

40. Using a table of standard electrode potentials, decide which of the following statements is completely true.
    1. 1. Cu²⁺ can oxidize H₂, and Fe can reduce Mn²⁺.
       2. Ni²⁺ can oxidize Cu²⁺, and Fe²⁺ can reduce H⁺.
       3. Fe²⁺ can oxidize H₂, and Fe²⁺ can reduce Au³⁺.
       4. Br₂ can oxidize Ni, and H₂ can reduce Mn²⁺.
       5. H⁺ can oxidize Fe, and Ni can reduce Br₂.

                                                        

 

41. Using a table of standard reduction potentials, determine which of these reactions (if any) is/are nonspontaneous in the direction indicated at 25°C.
    1. 1. 2Fe³⁺ + 2Cl^– → 2Fe²⁺ + Cl₂(g)        D)        A and B
       2. 2Fe³⁺ + 2Br^– → 2Fe²⁺ + Br₂(l)         E)        All are spontaneous.
       3. 2Fe³⁺ + 2I^– → 2Fe²⁺ + I₂(s)              

                                                         

 

42. Which one of the following reagents is capable of transforming Cu²⁺(1 M) to Cu(s)?

            A)  I^– (1 M)    B)  Ni(s)    C)  Al³⁺ (1 M)    D)  F^– (1 M)    E)  Ag(s)

                                                        

 

43. Which one of the following reagents is capable of transforming Br^– (aq) to Br₂(l) under standard-state conditions?

            A)  I^– (aq)    B)  NO₃^– (aq)    C)  Ag⁺ (aq)    D)  Al³⁺ (aq)    E)  Au³⁺ (aq)

                                                        

 

44. Which one of the following reagents is capable of transforming Cu(s) to Cu²⁺ (1 M)?

  A)  I^– (1 M)    B)  Ni(s)    C)  Ag⁺ (1 M)    D)  Al³⁺ (1 M)    E)  H⁺ (1 M)         

 

45. Which one of the following reagents is capable of transforming Fe³⁺ (1 M) to Fe²⁺ (1 M)?

             A)  H₂(1 atm)    B)  NO₃^– (1 M)    C)  O₂(1 atm)    D)  Br^– (1 M)    E)  H⁺ (1 M)

                                                   

 

46. The half-cell reaction for the oxidation of H₂O(l) to O₂(g) is given below.

             2H₂O(l) → O₂(g) + 4H⁺(aq) + 4e^–

  Which choice lists all of the following species that can oxidize H₂O to O₂(g) under standard-state conditions?

             MnO₄^–(aq), Cl₂(g), Pb²⁺(aq), Cl^– (aq), Ag⁺(aq)

1. 1. 1. Cl^–(aq) only             D)        Cl^–(aq) and MnO₄^–(aq)
      2. Cl₂(g) only   E)        MnO₄^–(aq) and Cl₂(g)
      3. Pb²⁺(aq) and Ag⁺(aq)                       

                                                        

 

47. Which of the following species is the strongest oxidizing agent under standard-state conditions?

            A)  Ag⁺(aq)    B)  H₂(g)    C)  H⁺(aq)    D)  Cl₂(g)    E)  Al³⁺(aq)

                                                         

 

48. Consider the following reaction:  2Fe²⁺(aq) + Cu²⁺ → 2Fe³⁺(aq) + Cu.

             When the reaction comes to equilibrium, what is the cell voltage?

            A)  0.43 V    B)  1.11 V    C)  0.78 V    D)  –0.43 V    E)  0 V

                       Category:  Easy     Section:  19.4

 

49. Determine the equilibrium constant, K_eq, at 25°C for the reaction

            2Br^– (aq) + I₂(s)         Br₂(l) + 2I^– (aq)

  A)  5.7 × 10^–19    B)  18.30    C)  1.7 × 10⁵⁴    D)  1.9 × 10¹⁸    E)  5.7 × 10^–55  

 

50. Determine the equilibrium constant (K_eq) at 25°C for the reaction

            Cl₂(g) + 2Br^– (aq)         2Cl^– (aq) + Br₂(l)

           A)  1.5 × 10^–10    B)  6.3 × 10⁹    C)  1.3 × 10⁴¹    D)  8.1 × 10⁴    E)  9.8

                                                        

 

51. Given the following standard reduction potentials,

                  Ag⁺(aq) + e^– → Ag(s)                                   E° = 0.80 V

                  AgCN(s) + e^– → Ag(s) + CN^–(aq)               E° = –0.01 V

             calculate the solubility product of AgCN at 25°C.

1. 1. 1. 4.3 × 10^–14   D)        5.1 × 10¹³
      2. 2.3 × 10¹³    E)        None of these
      3. 2.1 × 10^–14                

                                                        

 

52. Given the following standard reduction potentials,

                  Ag⁺(aq) + e^– → Ag(s)                                      E° = 0.80 V

                  Ag(NH₃)₂⁺(aq) + e^– → Ag(s) + 2NH₃(aq)       E° = 0.04 V

             calculate the formation constant of Ag(NH₃)₂⁺ at 25°C.

1. 1. 1. 6.1 × 10^–15   D)        1.6 × 10¹⁴
      2. 1.5 × 10^–13   E)        None of these
      3. 6.9 × 10¹²                 

                                                        

 

53. For the electrochemical cell Ni(s) | Ni²⁺(1 M) || H⁺(1 M) | H₂(1 atm) | Pt(s), which one of the following changes will cause a decrease in the cell voltage?
    1. 1. Increase the pressure of H₂ to 2.0 atm .
       2. Decrease the mass of the nickel electrode.
       3. Lower the pH of the cell electrolyte.
       4. Decrease the concentration of Ni²⁺ ion.
       5. None of the above.

                                                   

 

54. For the electrochemical cell Pt(s) | H₂(1 atm) | H⁺(1 M) || Cu²⁺(1 M) | Cu(s), which one of the following changes will cause an increase in the cell voltage?
    1. 1. Lower the H₂(g) pressure.
       2. Increase the size/mass of the copper electrode.
       3. Lower the H⁺(aq) concentration.
       4. Decrease the concentration of Cu²⁺ ion.
       5. None of the above.

                                                        

 

55. Consider an electrochemical cell with the following cell reaction where all reactants and products are at standard-state conditions:  Cu²⁺(aq) + H₂(g)  →  Cu(s) + 2H⁺(aq). Predict the effect on the emf of this cell of adding NaOH solution to the hydrogen half-cell until the pH equals 7.0.
    1. 1. The emf will increase.
       2. The emf will decrease.
       3. No change in the emf will be observed.

                                                   

 

56. Calculate the cell emf for the following reaction at 25°C:   2Ag⁺(0.010 M) + H₂(1 atm) → 2Ag(s) + 2H⁺(pH = 10.0)

            A)  1.04 V    B)  1.27 V    C)  0.92 V    D)  0.56 V    E)  0.80 V

                                                        

 

57. Calculate the cell emf for the following reaction at 25°C:

             Ni(s) + 2Cu²⁺(0.010 M) → Ni²⁺(0.0010 M) + 2Cu⁺(1.0 M)

            A)  0.40 V    B)  –0.43 V    C)  0.43 V    D)  0.34 V    E)  0.37 V

                                                        

 

58. Calculate the cell emf for the following reaction:

             Cu²⁺(0.10 M) + H₂(1 atm) → Cu(s) + 2H⁺(pH = 3.00)

            A)  0.49 V    B)  0.19 V    C)  0.15 V    D)  0.40 V    E)  –0.34 V

                                                   

 

59. Calculate the cell voltage for the following reaction:

             Cu²⁺ (0.010 M) + H₂(1 atm) → Cu(s) + 2H⁺( pH = 7.0)

            A)  0.19 V    B)  –0.01 V    C)  0.34 V    D)  0.69 V    E)  0.49 V

                                                         

 

60. Consider an electrochemical cell based on the spontaneous reaction

             2AgCl(s) + Zn(s) → 2Ag(s) + 2Cl^–  + Zn²⁺.

             If the zinc ion concentration is kept constant at 1 M, and the chlorine ion concentration is decreased from 1 M to 0.001 M, the cell voltage should

1. 1. 1. increase by 0.06 V. D)        decrease by 0.18 V.
      2. increase by 0.18 V. E)        increase by 0.35 V.
      3. decrease by 0.06 V.                          

                                                        

 

61. Consider an electrochemical cell involving the overall reaction

             2AgBr(s) + Pb(s) → Pb²⁺ + 2Ag(s) + 2Br^–

                 Each half-reaction is carried out in a separate compartment.  The anion included in the lead half-cell is NO₃^–.  The cation in the silver half-cell is K⁺.  The two half-cells are connected by a KNO₃ salt bridge.  If [Pb²⁺] = 1.0 M, what concentration of Br^– ion will produce a cell emf of 0.25 V at 298 K?

            Given:  AgBr(s) + e^– → Ag + Br^–, E° = +0.07 V.

            A)  0.02 M    B)  0.14 M    C)  0.38 M    D)  1.0 M    E)  7.0 M

                                                        

 

62. The measured voltage of the cell Pt(s) | H₂ (1.0 atm) | H⁺(aq) || Ag⁺(1.0 M) | Ag(s) is 1.02 V at 25°C.  Calculate the pH of the solution.

            A)  1.86    B)  1.69    C)  3.72    D)  3.89    E)  7.43

                                                        

 

63. The measured voltage of a cell in which the following reaction occurs is 0.96 V:

  H₂(g, 1.0 atm) + 2Ag⁺(aq, 1.0 M) → 2H⁺(aq, pH = ?) + 2Ag(s)   Calculate the pH of the H⁺(aq) solution.

            A)  1.4    B)  2.7    C)  5.4    D)  7.1    E)  14.9

                                                        

 

64. If the measured voltage of the cell  Zn(s) | Zn²⁺(aq) || Ag⁺(aq) | Ag(s) is 1.37 V when the concentration of Zn²⁺ ion is 0.010 M, what is the Ag⁺ ion concentration?
    1. 1. 2.5 M           D)        2.6 × 10^–51 M
       2. 4.0 × 10^–9 M            E)        6.2 × 10^–5 M
       3. 6.2 × 10^–3 M                         

                                                        

 

65. Which one of the following reactions must be carried out in an electrolytic cell rather than in a galvanic cell?
    1. 1. Zn²⁺ + Ca → Zn + Ca²⁺             C)        2Al + 3Fe²⁺ → 2Al³⁺ + 3Fe
       2. Al³⁺ + 3Br^– → Al + (3/2)Br₂      D)        H₂ + I₂(s) → 2H⁺ + 2I^–

       

66. Predict the products obtained from electrolysis of a 1 M AlBr₃ solution.

            Note that 2H₂O(l) + 2e^– → H₂(g) + 2OH^–(aq), E°_red = –0.83 V, and

             O₂(g) + 4H⁺(aq) + 4e^– → 2H₂O(l), E°_red = +1.23 V

           A)  Al and Br₂    B)  Al and O₂    C)  H₂ and O₂    D)  H₂ and Br₂    E)  Al and H₂

                                                       

 

67. Predict the products of the electrolysis of aqueous aluminum bromide AlBr₃(aq).  (Balancing is not required.)

            A)  Al + Br₂    B)  Al + O₂ + H⁺    C)  H₂ + OH^–  + Br₂    D)  H₂ + O₂

                                                       

 

68. Calculate the minimum voltage required for the electrolysis of 1.0 M NaCl in neutral solution.

             2H₂O + 2Cl^– (1.0 M) → H₂(1 atm) + Cl₂(1 atm) + 2OH^– (1 × 10^–7 M)

            A)  2.19 V    B)  1.78 V    C)  0.41 V    D)  –0.41 V    E)  –1.78 V

                                                        

 

69. When an aqueous solution of AgNO₃ is electrolyzed, a gas is observed to form at the anode.  The gas is

            A)  H₂.    B)  O₂.    C)  NO.    D)  NO₂.

                                                       

 

70. The half-reaction that occurs at the cathode during electrolysis of an aqueous sodium iodide solution is
    1. 1. Na⁺ + e^– → Na.        D)        I₂ + 2e^– → 2I^–.
       2. Na → Na⁺ + e^–.        E)        2I^– → I₂ + 2e^–.
       3. 2H₂O + 2e^– → H₂ + 2OH^–.              

                                                       

 

71. The half-reaction that occurs at the cathode during electrolysis of an aqueous CuCl₂ solution is

            A)        Cu⁺ + e^– → Cu.           D)        Cl₂ + 2e^– → 2Cl^–.        B)        Cu²⁺ + e^– → Cu⁺.             E)        2Cl^– → Cl₂ + 2e^–.

            C)      2H₂O + 2e^– → H₂ + 2OH^–.                              

                                                       

 

72. The half-reaction occurring at the cathode during electrolysis of an aqueous copper(II) iodide solution is
    1. 1. I₂ + 2e^– → 2I^–.   D)        2I^–  → I₂ + 2e^–.
       2. Cu → Cu²⁺ + 2e^–.          E)        2e^– + 2H₂O → H₂ + 2OH^–.
       3. Cu²⁺ + 2e^– → Cu.                       

                                                       

 

73. Under standard-state conditions, which of the following half-reactions occurs at the cathode during the electrolysis of aqueous nickel sulfate at 25°C?

            A)        2H₂O → O₂ + 4H⁺ + 4e^–         C)        2H₂O + 2e^– → H₂ + 2OH^–      B)        Ni²⁺ + 2e^– → Ni    D)        Ni → Ni²⁺ + 2e^–

                                                       

 

74. The half-reaction that should occur at the anode during electrolysis of an aqueous potassium bromide solution is
    1. 1. Br₂ + 2e^– → 2Br^–.    D)        2Br^– → Br₂ + 2e^–.
       2. Na → Na⁺ + e^–.        E)        2H₂O → O₂ + 4H⁺ + 4e^–.
       3. Na⁺ + e^– → Na.                     

                                                       

 

75. How many coulombs of charge are required to cause reduction of 0.20 mole of Cr³⁺ to Cr?

  A)  0.60 C    B)  3.0 C    C)  2.9 × 10⁴ C    D)  5.8 × 10⁴ C    E)  9.65 × 10⁴ C        Category:  Easy     Section:  19.8

 

76. How many coulombs of charge are required to cause reduction of 0.25 mole of Cu²⁺ to Cu?

  A)  0.25 C    B)  0.50 C    C)  1.2 × 10⁴ C    D)  2.4 × 10⁴ C    E)  4.8 × 10⁴ C        Category:  Easy     Section:  19.8

 

77. How many faradays are transferred in an electrolytic cell when a current of 2.0 amperes flows for 12 hours?

            A)  24 F    B)  8.6 × 10⁴ F    C)  0.90 F    D)  6.2 × 10 ^–3 F    E)  1.1 F

                                                        

 

78. A metal object is to be gold-plated by an electrolytic procedure using aqueous AuCl₃ electrolyte.  Calculate the number of moles of gold deposited in 3.0 min by a constant current of 10. A.
    1. 1. 6.2 × 10^–3 mol         D)        3.5 × 10^–5 mol
       2. 9.3 × 10^–3 mol         E)        160 mol
       3. 1.8 × 10^–2 mol                      

                                                   

 

79. A current of 0.80 A was applied to an electrolytic cell containing molten CdCl₂ for 2.5 hours. Calculate the mass of cadmium metal deposited.

 A)  3.2 × 10^–7 g    B)  1.2 × 10^–3 g    C)  4.2 g    D)  8.4 g    E)  16.8 g

       

80. A current of 2.50 A was passed through an electrolytic cell containing molten CaCl₂ for

4.50 hours.  How many moles of calcium metal should be deposited?

1. 1. 1. 5.83 × 10^–5 mol       D)        0.840 mol
      2. 0.210 mol    E)        1.95 × 10⁹ mol
      3. 0.420 mol                 

                                                        

 

81. How many grams of nickel would be electroplated by passing a constant current of 7.2 A through a solution of NiSO₄ for 90.0 min?

            A)  0.20 g    B)  0.40 g    C)  12 g    D)  24 g    E)  47 g

                                                        

 

82. How many coulombs (C) of electrical charge must pass through an electrolytic cell to reduce 0.44 mol Ca²⁺ ion to calcium metal?

            A)  190,000 C    B)  85,000 C    C)  21,000 C    D)  42,500 C    E)  0.88 C

                                                        

 

83. How many grams of chromium would be electroplated by passing a constant current of

5.2 amperes through a solution containing CrCl₃ for 45.0 min?

            A)  9.3 × 10^–4 g    B)  0.042 g    C)  2.5 g    D)  24 g    E)  2.3 × 10¹⁰ g

                                                        

 

84. How many coulombs would be required to electroplate 35.0 grams of chromium by passing an electrical current through a solution containing CrCl₃?
    1. 1. 6.50 × 10⁴ C            D)        1.95 × 10⁵ C
       2. 2.16 × 10⁴ C            E)        1.01 × 10⁷ C
       3. 6.40 × 10⁴ C                         

                                                         

 

85. How many minutes would be required to electroplate 25.0 grams of chromium by passing a constant current of 4.80 amperes through a solution containing CrCl₃?

             A)  483 min    B)  161 min    C)  322 min    D)  2.01 × 10⁴ min    E)  1.11 × 10⁴ min           

 

86. How long will it take to produce 78 g of Al metal by the reduction of Al³⁺ in an electrolytic cell with a current of 2.0 A?

            A)  0.01 s    B)  420 s    C)  13 h    D)  116 h    E)  1.0 × 10¹² s

 

87. Aluminum does not corrode as does iron, because A)         Al does not react with O₂.
    1. 2. A protective layer of Al₂O₃ forms on the metal surface.
       3. Al is harder to oxidize than is Fe.
       4. Fe gives cathodic protection to Al.
       5. The electrical circuit cannot be completed on an Al surface.

 

88. Iron objects such as storage tanks and underground pipelines can be protected from corrosion by connecting them through a wire to a piece of

            A)  Pb    B)  Ag    C)  Sn    D)  Mg    E)  Cu

 

89. Which element is associated with the term "galvanized"?

            A)  Ga    B)  Zn    C)  Cd    D)  Hg    E)  Pb

 

90. Find the emf of the cell described by the cell diagram

            Fe | Fe²⁺ (1.500M) || Au³⁺ (0.00400M) | Au

            A)  1.99 V    B)  1.89 V    C)  1.94 V    D)  1.66 V    E)  1.91 V

 

91. Find the emf of the cell described by the cell diagram

             Ni | Ni²⁺ (0.750 M) || Cu²⁺ (0.0500 M) | Cu

            A)  0.62 V    B)  0.52 V    C)  0.59 V    D)  0.66 V    E)  0.56 V

 

92. Consider the reaction Pb(s) + 2H⁺(aq) → Pb²⁺(aq) + H₂(g).  If the hydrogen gas pressure is maintained at 1.00 atm and the lead(II) ion concentration is 0.025 M, find the pH at which this reaction is at equilibrium.

            A)  2.8    B)  3.2    C)  6.0    D)  3.0    E)  5.7

 

93. Suppose the reaction Pb(s) + 2H⁺(aq) → Pb²⁺(aq) + H₂(g) is carried out at pH = 4.00 and at a hydrogen gas pressure of 1.00 atm.  The concentration of lead(II) ions that causes this reaction to be at equilibrium is

  A)  2.5 M    B)  1.6 × 10^–2 M    C)  2.5 × 10^–4 M    D)  1.6 × 10^–6 M    E)  0.40 M  

 

94. Which of these metals will be oxidized in hydrochloric acid under standard conditions at 25°C?

            A)  Ag    B)  Au    C)  Hg    D)  Cu    E)  Zn

 

95. Which of these metals will not be oxidized in hydrochloric acid under standard conditions at 25°C?

 A)  Al    B)  Fe    C)  Ag    D)  Ni    E)  Mg

96. Which of these metals will reduce water to hydrogen in basic solution under standard conditions?

            A)  Pb    B)  Fe    C)  Zn    D)  Na    E)  Ag

 

97. Which of these metals will not reduce water to hydrogen in basic solution under standard conditions?

            A)  Cd    B)  Sr    C)  Mg    D)  Ba    E)  K

 

98. Consider the reaction 2Fe³⁺(aq) + Fe(s)  3Fe²⁺(aq).  Find the equilibrium constant for this reaction at 25°C.

  A)  1 × 10¹¹    B)  8 × 10⁴⁰    C)  3 × 10²⁰    D)  7 × 10^–12    E)  1 × 10^–41  

 

99. Consider the reaction Hg²⁺(aq) + Hg(l)         Hg₂²⁺(aq).  Find the equilibrium constant for this reaction at 25°C.

  A)  2 × 10¹    B)  6 × 10⁵⁹    C)  4 × 10^–3    D)  7 × 10^–2    E)  2 × 10²  

 

100. Complete and balance the following redox reaction under acidic conditions:

             Cr₂O₇^2-(aq) + I^-(aq) → Cr³⁺(aq) + IO₃^-(aq)

                                          

 

101. Complete and balance the following redox reaction under acidic conditions:

            Cu(s) + NO₃^-(aq) → Cu²⁺(aq) + NO₂(g)

                                          

 

102. Complete and balance the following redox reaction under acidic conditions:

             ClO₂^-(aq) → ClO₂(g) + Cl^-(aq)

                                          

 

103. Complete and balance the following redox reaction under acidic conditions:

             MnO₄^-(aq) + C₂O₄^2-(aq) → Mn²⁺(aq) + CO₂(aq)

                                          

 

104. Complete and balance the following redox reaction under basic conditions:

 PO₃^3-(aq) + MnO₄^-(aq) → PO₄^3-(aq) + MnO₂(s)

105. Complete and balance the following redox reaction under basic conditions:

             CrO₄^2-(aq) + SO₃^2-(aq) → Cr(OH)₃(s) + SO₄^2-(aq)

                                          

 

106. Gold can be electrochemically "plated" onto a less expensive metal by submerging the object in a solution of gold(III) and applying an electric current to drive the following half-reaction:

             Au³⁺ + 3e^– → Au(s)

             How many grams of gold will be formed if a current of 2.37 A is applied for 71 minutes?

                                           

 

107. What current is needed to deposit 0.500 g of chromium metal from a solution of Cr³⁺ in a period of 1.00 hr?

                                           

 

108. Will H₂(g) form when Fe is placed in 1.0 M HCl?

                                           

 

109. Will H₂(g) form when Ag is placed in 1.0 M HCl?

                                           

 

110. Will H₂(g) form when Cu is placed in 1.0 M HCl?

                                           

 

111. Will H₂(g) form when Sn is placed in 1.0 M HCl?

                                           

 

112. When a solution of a certain gadolinium salt is electrolyzed with a current of 1.0 A for 2.0 h, 0.025 mol of Gd metal forms.  Calculate the charge on the gadolinium ion in the salt.

                                          

 

113. Given the following standard reduction potentials in acid solution

                  O₂ + 4H⁺ + 4e^– → 2H₂O              E° = +1.23 V

Sn⁴⁺ + 2e^– → Sn²⁺ E° = +0.13 V Zn²⁺ + 2e^– → Zn(s) E° = –0.76 V

            write the formula of the weakest reducing agent.

                                           

 

114. Given the following standard reduction potentials in acid solution

                  O₂ + 4H⁺ + 4e^– → 2H₂O              E° = +1.23 V

                  Sn⁴⁺ + 2e^– → Sn²⁺                        E° = +0.13 V

                  Zn²⁺ + 2e^– → Zn(s)                      E° = –0.76 V

             write the formula of the strongest reducing agent.

                                           

 

115. Given the following standard reduction potentials in acid solution

                  O₂ + 4H⁺ + 4e^– → 2H₂O              E° = +1.23 V

                  Sn⁴⁺ + 2e^– → Sn²⁺                        E° = +0.13 V

                  Zn²⁺ + 2e^– → Zn(s)                      E° = –0.76 V

             write the formula of the strongest oxidizing agent.

                                           

 

116. Given the following standard reduction potentials in acid solution

                  O₂ + 4H⁺ + 4e^– → 2H₂O              E° = +1.23 V

                  Sn⁴⁺ + 2e^– → Sn²⁺                        E° = +0.13 V

                  Zn²⁺ + 2e^– → Zn(s)                      E° = –0.76 V

 

  write a balanced equation for a spontaneous reaction which involves the tin and zinc redox couples.

                                           

 

117. Determine the equilibrium constant for the following reaction at 25°C.

  2I^– (aq) + Br₂(l)  I₂(s) + 2Br^–(aq).  Ans: 1.8 × 10¹⁸

                                           

 

118. Determine the equilibrium constant for the following reaction at 298 K.  2Fe³⁺(aq) + H₂O₂(aq)  2H⁺(aq) + O₂(g) + 2Fe²⁺(aq).

                                           

 

119. If the cell emf of a Zn–Cu cell is 0.80 V when the concentration of Zn²⁺ is 2.0 M, what is the concentration of Cu²⁺?

 

120. Aluminum metal is formed by the electrolysis of Al₂O₃ in molten cryolite.  How many grams of Al are produced when 6.50 × 10³ C pass through the cell?

                                          

 

121. Aluminum metal is formed by the electrolysis of Al₂O₃ in molten cryolite.  How many minutes are required to form 10.0 g of Al using a current of 30 A?

                                          

 

122. How many moles of H₂ are produced by 5.00 A of current passing through a cell containing aqueous NaCl for 4.00 × 10² s?

                                          

 

123. What concentration of Ni²⁺ ion remains in solution after electrolysis of 100. mL of 0.250 M NiSO₄ solution when using a current of 2.40 amperes for 30.0 minutes?  Assume Ni metal is plated out.

                                          

 

124. Calculate the cell emf for the following reaction at 25°C:

            2Ag⁺(0.010 M) + H₂(1 atm) → 2Ag(s) + 2H⁺(pH = 6.0)

                                           

 

125. A standard hydrogen electrode is immersed in an acetic acid solution.  This electrode is connected by an external circuit to an iron nail dipping into 0.10 M FeCl₂.  If E_cell is found to be 0.24 V, what is the pH of the acetic acid solution?

 

126. Consider the reaction Fe + Sn²⁺(1 × 10^–3 M) → Fe²⁺(1.0 M) + Sn.  Calculate the voltage theoretically generated by the cell.

                                           

 

127. Consider the reaction Fe + Sn²⁺(1 × 10^–3 M) → Fe²⁺(1.0 M) + Sn.  Sketch a diagram of such a cell.  Be sure to label the anode and cathode, and show the direction of electron flow in the external circuit.  Write the half-reactions occurring at the electrodes.

 

128. Consider the reaction Fe + Sn²⁺(1 × 10^–3 M) → Fe²⁺(1.0 M) + Sn.  In what direction do the Cl^–  ions in the KCl salt bridge flow?

 

129. How many moles of silver metal are produced in 1.2 hours by the electrolysis of AgNO₃(aq) using a current of 6.0 A?

                                           

 

130. How many grams of copper are deposited on the cathode of an electrolytic cell if an electric current of 2.00 A is passed through a solution of CuSO₄ for a period of 19.0 min?

                                           

 

131. An electroplating solution is made up of nickel(II) sulfate. How much time would it take to deposit 0.500 g of metallic nickel on a custom car part using a current of 3.00 A?

                                           

 

132. Many different ways have been proposed to make batteries. One cell is set up with copper and lead electrodes in contact with CuSO₄(aq) and Pb(NO₃)₂ (aq), respectively. If the Pb²⁺ and Cu²⁺ concentrations are each 1.0 M, what is the overall cell potential?

                  Pb²⁺ + 2e^– → Pb                         E° = –0.22 V

                  Cu²⁺ + 2e^– → Cu                        E° = +0.34 V

 

                                       

 

133. A battery is constructed by placing copper and lead electrodes in contact with 1.0 molar CuSO₄(aq) and Pb(NO₃)₂ (aq) solutions, respectively.  If sulfuric acid is then added to the Pb(NO₃)₂ solution, thereby forming a precipitate of PbSO₄, what will happen to the cell potential?

                                           

 

134. You wish to electroplate a metal utensil with a surface area of 736 cm² with gold to give an average thickness of 0.025 mm over the entire surface. Starting with a solution of excess Au³⁺ and applying a constant current of 14 A, how long will it take to electroplate the utensil? The density of gold is 19.3 g/cm³.

                                          

 

135. In an electrolytic cell, electrical energy is used to cause a chemical reaction to occur that would otherwise be nonspontaneous.

 

136. When an aqueous solution of NaCl is electrolyzed, Na(l) is produced at the cathode, and Cl₂(g) is evolved at the anode.

 

137. The reduction of 1.00 mole of Cr³⁺ to Cr requires 9.65 × 10⁴ C of electrical charge.

 

138. The electrochemical cell that utilizes the reaction Zn + Cu²⁺ (1 M) → Zn²⁺ (1 M) + Cu will have a lower cell emf when the concentrations of Zn²⁺ and Cu²⁺ ions are decreased to 0.1 M.