Electrochemistry- Problem Set

Electrochemistry- Problem Set

Electrochemistry- Problem Set

l. For the reaction: 2 Fe + 3 CdCl2 <=====> 2 FeCl3 + 3 Cd, which statement is true?

a) Fe is the oxidizing agentb) Cd undergoes oxidation

c) Cd is the reducing agentd) Fe undergoes oxidation

2. What is the potential in volts for the spontaneous reaction between the Ag/Ag+ and Zn/Zn+2 half-


Zn2+ + 2 e-  Zn (E° = -0.763 V); Ag+ + e-  Ag (E°= 0.799 V)

a) -2.361 V b) -1.562 V c) 1.562 V d) 2.361 V

3. What is the oxidation state of arsenic in H3AsO3?

a) +1 b) +3 c) +5 d) +7

4. When the reaction below is balanced, what is the coefficient of H+?

MnO4- + NO2- + H+ -----> MnO2 + NO3- + H2O

a) 2 b) 4 c) 6 d) 8

5. What is the [Cu2+] in the cell Zn / Zn2+ (0.05M) // Cu2+ (? M) / Cu if the cell voltage is 1.03V?

Zn2+ + 2 e-  Zn (E° = -0.763 V); Cu2+ + 2 e-  Cu (E° = 0.337 V)

a) 0.12 Mb) 0.0002 M c) 0.05 M d) 0.0035

6. 10 amps are passed through molten aluminum chloride for 5.5 hours. How many grams of aluminum

metal could be produced by this electrolysis?

a) 18.5 gb) 55.4 g c) 91.2 g d) 273 g

7. In the electroplating of silver from cyanide solution, the cathode reaction is:

Ag(CN)2- + 1 e-  Ag (s) + 2 CN-

How many grams of silver should be deposited by a current of 4.5 A in 28.0 minutes?

a) 0.141 gb) 4.23 g c) 8.46 g d) 12.53 g


8. Given the standard electrode potentials:

Ni2+ + 2 e-  Ni(s) (E ° = -0.23 V); Cr3+ + 3 e-  Cr(s) (E ° = -0.74 V)

Which pair of substances with react spontaneously?

a) Ni2+ with Cr3+ b) Ni with Cr3+ c) Ni2+ with Cr d) Ni with Cr

9. Aluminum oxide may be electrolyzed at 1000°C to furnish aluminum metal.

The cathode reaction is: Al3+ + 3 e-  Al

To prepare 5.12 kg of aluminum metal would require how many coulombs of electricity?

a) 5.49 x 107 Cb) 1.83 x 107 Cc) 5.49 x 104 Cd) 5.49 x 101 C

10. The electrolysis of molten sodium chloride is carried out in an electrochemical cell.

a) Write the balanced half-cell reactions that occur at each electrode, indicating whether each takes

place at the anode or cathode.

b) If a current of 1.06 amps is used, how long will it take to make 1.00 L of Cl2(g) at 24°C

and 742 mm Hg?

c) Calculate K and G° for this system.

11. A voltaic cell is designed using the following reaction: Zn + 2 Ag+  Zn+2 + 2 Ag

a) Determine the standard voltage, E°, for this reaction.

b) Suppose the concentration of Zn+2 in the Zn/Zn+2 half-cell is maintained at l.00 M. Excess

hydrochloric acid is added to the Ag/Ag+ half-cell, precipitating AgCl and making the concentration of Cl- = 0.100 M. Under these conditions, the cell voltage is found to be 1.04 V. Calculate the concentration of Ag+ in the Ag/Ag+ half-cell.

c) Use the information in (b) to calculate the Ksp of AgCl.

12. An electrochemical cell consists of a nickel electrode in an acidic solution of l.00-molar Ni(NO3)2

connected by a salt bridge to a second component with an aluminum electrode in an acidic solution of

l.00-molar AlCl3.

a) Write an equation for the half-cell reaction occurring at each electrode. Indicate whether each reaction

occurs at the anode or the cathode.

b) Write a net ionic equation for the overall spontaneous cell reaction that occurs when the circuit is

complete. Calculate the standard voltage, E°, for this reaction.

c) Calculate the change in voltage when the cell described above has initial concentrations of

0.500 molar Ni(NO3) 2 and 0.750 molar AlCl3.

13. E° = 1.101 volt at 25°C. for the reaction: M(s) + Cu2+(aq)  M2+(aq) + Cu (s)

a) Determine the standard electrode potential for the reduction half-reaction: M2+ + 2 e-  M

b) A cell is constructed in which the reaction above occurs. All substances are initially in their standard

states, and equal volumes of the solutions are used. The cell is then discharged. Calculate the value of the cell potential, E, when [Cu2+] has dropped to 0.20 molar.