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Homework answers / question archive / University of South Florida Course title: CHM 2046 1)Under standard conditions, which of the following operations results in a spontaneous chemical reaction taking place? A piece of aluminum metal is placed in an aqueous solution of potassium nitrate

University of South Florida Course title: CHM 2046 1)Under standard conditions, which of the following operations results in a spontaneous chemical reaction taking place? A piece of aluminum metal is placed in an aqueous solution of potassium nitrate

Chemistry

University of South Florida

Course title: CHM 2046

1)Under standard conditions, which of the following operations results in a spontaneous chemical reaction taking place?

    1. A piece of aluminum metal is placed in an aqueous solution of potassium nitrate.

 

    1. Iodine crystals are added to an aqueous solution of sodium chloride.
    2. A piece of silver metal is placed in an aqueous solutions of copper(II) nitrate.
    3. Chlorine gas is bubbled through an aqueous solution of sodium bromide.
    4. At least two of the above (A-D) result in a spontaneous chemical reaction.

 

 

               

 

  1.  

4

A common car battery consists of six identical cells, each of which carries out the reaction: Pb + PbO2 + 2HSO + 2H+ ® 2PbSO4 + 2H2O

The value of     for such a cell is 2.039 V. Calculate DG° at 25 °C for the reaction. A) –196.7 kJ

B)     –98.37 kJ

C)     –393.5 kJ

D)   –786.9 kJ

E)     –590.2 kJ

 

 

               

 

  1. What is DG° for the following electrochemical equation? (e°red(Ag+/Ag) = 0.800 V, e°red(Cd2+/Cd) = –0.403 V) 2Ag(s) + Cd2+(aq) ® 2Ag+(aq) + Cd(s)
    1. –232 kJ/mol
    2. 116 kJ/mol
    3. 232 kJ/mol
    4. 464 kJ/mol
    5. –464 kJ/mol

 

 

               

 

Consider an electrochemical cell with a zinc electrode immersed in 1.0 M Zn2+ and a nickel electrode immersed in

0.10 M Ni2+.

Zn2+ + 2e- ® Zn                         e° = –0.76 V

Ni2+ + 2e- ® Ni                          e° = –0.23 V

 

 

  1. Calculate the concentration of Ni2+ if the cell is allowed to run to equilibrium at 25°C. A) 1.10 M

B)     0.20 M

C)     0.10 M

  1. 0 M
  2. none of these

 

 

               

 

 

  1. Calculate e at 25°C for the cell shown below, given the following data:

 

 

Ni

Ag

 

 

 

 

 

Ksp for AgCl = 1.6 ´ 10–10 A) 0.83 V

B)     0.54 V

C)     1.01 V

D)   2.98 V

E)     cannot be determined from the data given

 

 

               

 

  1. The following question refers to the following system:

 

3                                                                                                                   2

3Ag(s) + NO (aq) + 4H+(aq) ® 3Ag+(aq) + NO(g) + 2H O(l)

 

3                                                                                                  2

Anode reaction:          Ag ® Ag+(aq) + 1e                                                                          e° = –0.7990 V Cathode reaction:      NO (aq) + 4H+(aq) + 3e ® NO(g) + 2H O(l)                      e° = 0.9636 V

Determine the equilibrium constant at 25°C.

A)   6.097 ×102

B)     2.965 ×1089

C)     4.412 ×10- 9

D)   2.266 ×108

E)     3.117

 

 

               

 

  1. A galvanic cell consists of a left compartment with a tin electrode in contact with 0.1 M Sn(NO3)2(aq) and a right compartment with a lead electrode in contact with 1 ´ 10–3 M Pb(NO3)2(aq). The relevant reduction potentials are:

Pb2+ + 2e ® Pb                  e° = –0.13 V

Sn2+ + 2e ® Sn                   e° = –0.14 V

When this cell is allowed to discharge spontaneously at 25°C, which of the following statements is true?

    1. Electrons will flow from left to right through the wire.
    2. Pb2+ ions will be reduced to Pb metal.
    3. The concentration of Sn2+ ions in the left compartment will increase.
    4. The tin electrode will be the cathode.
    5. No noticeable change will occur, because the cell is at equilibrium.

 

 

               

 

  1. Consider the galvanic cell shown below (the contents of each half-cell are written beneath each compartment):

 

 

Pt

Cr

 

 

0.50 M Br2                 0.20 M Cr3+

0.10 M Br

The standard reduction potentials are as follows:

Cr3+(aq) + 3e ® Cr(s)                        e° = –0.733 V

Br2(aq) + 2e ® 2Br(aq)                  e° = +1.090 V

What is the value of e for this cell at 25°C? A) 1.759 V

B)     1.823 V

C)     1.837 V

D)   1.887 V

E)     2.207 V

 

 

               

 

 

  1. Consider an electrochemical cell with a zinc electrode immersed in a solution of Zn2+ and a silver electrode immersed in a solution of Ag+.

Zn2+ + 2e ® Zn                         e° = –0.76 V

Ag+ + e ® Ag                              e° = 0.80 V

 

 

 

If

A)

is 0.050 M and

1.46 V

B)

1.76 V

C)

1.36 V

D)

1.66 V

E)

1.63 V

 

is 11.26 M, calculate e.

 

 

 

 

 

 

               

 

 

  1. Consider an electrochemical cell with a zinc electrode immersed in 1.0 M Zn2+ and a nickel electrode immersed in

0.39 M Ni2+.

Zn2+ + 2e ® Zn                   e° = –0.76 V

Ni2+ + 2e ® Ni                    e° = –0.23 V

Calculate e for this cell. A) 0.54 V

B)     0.52 V

C)     0.51 V

D)   0.53 V

E)     0.98 V

 

 

               

 

 

  1. A cell is set up with copper and lead electrodes in contact with CuSO4(aq) and Pb(NO3)2(aq), respectively, at 25°C. The standard reduction potentials are:

Pb2+ + 2e ® Pb                  e° = –0.13 V

Cu2+ + 2e ® Cu                  e° = +0.34 V

If sulfuric acid is added to the Pb(NO3)2 solution, forming a precipitate of PbSO4, the cell potential:

    1. increases
    2. decreases
    3. is unchanged
    4. can't tell what will happen
    5. none of these

 

 

 

 

  1. A concentration cell is constructed using two Ni electrodes with Ni2+ concentrations of 1.0 M and 1.00 ´ 10–4 M in the two half-cells. The reduction potential of Ni2+ is –0.23 V. Calculate the potential of the cell at 25°C.

A)   –0.368 V

B)     +0.132 V

C)     –0.132 V

D)   +0.118 V

E)     +0.0592 V

 

 

               

 

 

  1. The standard potential for the reaction A + B        C + D is 1.50 volts. The equilibrium constant K for this reaction at 25°C is:

A) 2.5 ´ 1025

B)     4.0 ´ 10–26

C)     25.4

D)   –25.4

E)     not enough information given

 

 

               

  1. The reduction potentials for Ni2+ and Sn2+ are as follows: Ni2+ + 2e ® Ni, e° = –0.231 V

Sn2+ + 2e ® Sn, e° = –0.140 V

Calculate the equilibrium constant at 25 °C for the reaction:

 

 
   
 

 

Sn2+ + Ni         Sn + Ni2+

A)  3.6 ×1012

B)     35

C)     5.9

D)  8.3 ×10- 4

E)     1.2 ×103

 

               

 

  1. An excess of finely divided iron is stirred up with a solution that contains Cu2+ ion, and the system is allowed to come to equilibrium. The solid materials are then filtered off and electrodes of solid copper and solid iron are inserted into the remaining solution. What is the value of the ratio [Fe2+]/[Cu2+] at 25°C?

 

The following standard reduction potentials apply:

Fe2+(aq) + 2e ® Fe(s)               e° = –0.44 V

Cu2+(aq) + 2e ® Cu(s)              e° = +0.34 V

 

    1. 1
    2. 0

C)     2.5 ´ 1026

D) 4.4 ´ 10–27

E)     none of these

 

 

               

 

  1. An excess of finely divided iron is stirred up with a solution that contains Cu2+ ion, and the system is allowed to come to equilibrium. The solid materials are then filtered off and electrodes of solid copper and solid iron are inserted into the remaining solution. What potential develops between these two electrodes at 25°C?
    1. 0

B)     –0.78 V

C)     0.592 V

D)   0.296 V

E)     not enough information given

 

 

 

 

  1. The equilibrium constant at 25°C for the reaction Al + 3Cu2+           2Al3+ + 3Cu is approximately A) 10203

B)     1034

C)     1068

D)   10–203

E)     none of these

 

 

 

 

  1. A concentration cell is constructed with copper electrodes and Cu2+ in each compartment. In one compartment, the [Cu2+] = 1.0 ´ 10–3 M and in the other compartment, the [Cu2+] = 2.0 M. Calculate the potential for this cell at 25°C. The standard reduction potential for Cu2+ is +0.34 V.

A) 0.44 V

B)     –0.44 V

C)     0.098 V

D) –0.098 V

E)     0.78 V

 

 

               

 

 

  1. Calculate the solubility product of silver iodide at 25°C given the following data:

e°(V)

AgI(s) + e ® Ag(s) + I                     –0.15

I2(s) + 2e ® 2I                                   +0.54

Ag+ + e ® Ag(s)                                     +0.80

 

A)   2.9 ´ 10–3

B)     1.9 ´ 10–4

C)     2.1 ´ 10–12

D)   8.4 ´ 10–17

E)     6.1 ´ 10–26

 

               

 

  1. Using the following data to calculate Ksp for PbSO4.

 

4

PbO2 + 4H+ + SO 2– + 2e ® PbSO4(s) + 2H2O                        +1.69

PbO2 + 4H+ + 2e ® Pb2+ + 2H2O                                                 +1.46

 

A) 4.0 ´ 10106

B)     2.5 ´ 10–107

C)     5.9 ´ 107

D) 1.7 ´ 10–8

E)     None of these is within 5% of the correct answer.

 

 

               

 

  1. In which of the following cases must e be equal to zero?
    1. In any cell at equilibrium.
    2. In a concentration cell.
    3. e can never be equal to zero.
    4. Choices A and B are both correct.
    5. None of these.

 

 

 

 

  1. Which of the following statements is/are correct?
    1. The value of e° is equal to zero in a concentration cell, and the value of e is equal to zero in any cell at equilibrium.
    2. The value of e° can be equal to zero in a concentration cell, and the value of e must be equal to zero in a concentration cell.
    3. The values of e° and e are equal to zero in any cell at equilibrium.
    4. e° can never be equal to zero.
    5. At least two of the above choice (A-D) are correct.

 

 

 

 

  1. Which of the following statements is true about a voltaic cell for which e°cell = 1.00 V?
    1. It has DG° > 0.
    2. The system is at equilibrium.
    3. It has K = 1.
    4. The cathode is at a higher energy than the anode.
    5. The reaction is spontaneous.

 

 

               

 

  1. If a reducing agent M reacts with an oxidizing agent N+ to give M+ and N, and the equilibrium constant for the reaction is 1.6, then what is the e° value for the oxidation-reduction reaction?

A) 0.012 V

B)     –0.012 V

C)     0.0060 V

D) –0.0060 V

E)     0.024 V

 

 

 

 

  1. If you could increase the concentration of Zn2+, which of the following is true about the cell potential?
    1. It would increase.
    2. It would decrease.
    3. It would remain constant.
    4. Cannot be determined.
    5. None of these (A-D).

 

 

 

 

  1. If you could increase the concentration of Al3+, which of the following is true about the cell potential?
    1. It would increase.
    2. It would decrease.
    3. It would remain constant.
    4. Cannot be determined.
    5. None of these (A-D).

 

 

               

 

  1. A concentration cell is constructed using two metal (M) electrodes with M2+ concentrations of 0.10 M and 1.00 ´ 10–5 M in the two half-cells. Determine the reduction potential of M2+ given that the potential of the cell at 25°C is 0.118 V.
    1. 0 V

B)     +0.118 V

C)     –0.118 V

  1. Cannot be determined with the information given.
  2. None of the above.

 

 

               

 

 

  1. What is e of the following cell reaction at 25°C? e°cell = 0.460 V. Cu(s) | Cu2+(0.014 M) || Ag+(0.20 M) | Ag(s)

A)   0.282 V

B)     0.491 V

C)     0.460 V

D)   0.473 V

E)     0.494 V

 

 

               

 

 

  1. For the cell Cu(s) | Cu2+ || Ag+ | Ag(s), the standard cell potential is 0.46 V. A cell using these reagents was made, and the observed potential was 0.16 V at 25oC. What is a possible explanation for the observed voltage?
    1. The Ag+ concentration was larger than the Cu2+ concentration.
    2. The Cu2+ concentration was larger than the Ag+ concentration.
    3. The Ag electrode was twice as large as the Cu electrode.
    4. The volume of the Cu2+ solution was larger than the volume of the Ag+ solution.
    5. The volume of the Ag+ solution was larger than the volume of the Cu2+ solution.

 

               

 

  1. What is the value of the reaction quotient, Q, for the voltaic cell constructed from the following two half-reactions when the Zn2+ concentration is 0.0103 M and the Ag+ concentration is 1.35 M?

Zn2+(aq) + 2e ® Zn(s); e° = –0.76 V

Ag+(aq) + e ® Ag(s); e° = 0.80 V

A)   177

B)     131

C)     1.25 ×10- 2

D)   7.63 ×10- 3

E)     5.65 ×10- 3

 

               

 

 

  1. In order to determine the identity of a particular lanthanide metal Sm, a voltaic cell is constructed at 25°C with the anode consisting of the lanthanide metal as the electrode immersed in a solution of 0.0819 M SmCl3, and the cathode consisting of a copper electrode immersed in a 1.00 M Cu(NO3)2 solution. The two half-reactions are as follows:

Sm(s) ® Sm3+(aq) + 3e Cu2+(aq) + 2e ® Cu(s)

The potential measured across the cell is 2.67 V. What is the identity of the metal?

 

Reduction Half-Reaction

 

e° (V)

 

Cu2+(aq) + 2e ® Cu(s) Ce3+(aq) + 3e ® Ce(s)

Dy3+(aq) + 3e ® Dy(s)

 

0.34

–2.336

–2.295

 

 

 

 

Eu3+(aq) + 3e ® Eu(s) Gd3+(aq) + 3e ® Gd(s) Sm3+(aq) + 3e ® Sm(s)

 

–1.991

–2.279

–2.304

 

 

 

 

A) Ce

 

 

 

 

 

 

B)     Eu

 

 

 

 

 

 

C)     Dy

 

 

 

 

 

 

D) Sm

 

 

 

 

 

 

E)     Gd

 

 

 

 

 

 

               

 

 

 

 

  1. Which of the following statements is true concerning the electrochemical cell described below at 25oC? Cu | Cu2+(0.816 M) || Cu2+(0.843] M) | Cu

Cu2+(aq) + 2e ® Cu(s); e° = 0.34 V

    1. The cell reaction is spontaneous with a cell potential of 4.81 mV.
    2. The cell reaction is spontaneous with a cell potential of 0.418 mV.
    3. The cell reaction is nonspontaneous with a cell potential of –0.418 mV.
    4. The cell reaction is nonspontaneous with a cell potential of –4.81 mV.
    5. The cell reaction is spontaneous with a cell potential of 0.340 V.

 

 

               

 

 

  1. What is the potential at 25°C for the following cell? Cr | Cr3+(0.015 M) || Ag+(0.00025 M) | Ag

 

Cr3+ + 3e–     Cr                e° = –0.73 V

Ag+ + e–      Ag                  e° = 0.80 V

A)   2.09 V

B)     1.35 V

C)     0.95 V

D)   1.71 V

E)     1.49 V

 

 

               

 

 

  1. Concentration cells work because standard reduction potentials are dependent on concentration.

 

ANS:      F                            

 

  1. Consider the hydrogen–oxygen fuel cell where:

 

 
   
 

 

H2(g) + O2(g)       H2O(l)          DG° = –237.18 kJ/mol H2

Which of the following statements is true?

    1. At standard conditions, the maximum work the fuel cell could do on the surroundings is 237.18 kJ/mol.
    2. In the real world, the actual amount of useful work the cell can do is less than 237.18 kJ.
    3. More energy is dissipated as waste heat in the fuel cell than in the reversible pathway.
    4. A, B, and C are all true.
    5. A, B, and C are all false.

 

 

               

 

  1. Which type of battery has been designed for use in space vehicles?
    1. lead storage
    2. alkaline dry cell
    3. mercury cells
    4. fuel cells
    5. silver cells

 

 

               

 

  1. Which of the following statements about batteries is false?
    1. A battery is a group of galvanic cells connected in series.
    2. Lead storage batteries contain lead at the anode and lead coated with lead dioxide at the cathode.
    3. The alkaline dry cell battery can last longer than a nickel-cadmium battery.
    4. A fuel cell is a galvanic cell for which the reactants are continuously supplied.
    5. Dry cell batteries are used in tape players and portable radios.

 

 

               

 

  1. Which of the following statements about corrosion is false?
    1. Patina is the layer of tarnish that gives silver a richer appearance.
    2. The oxidation of most metals by oxygen is spontaneous.
    3. Most metals will develop a thin oxide coating, which protects their internal atoms from oxidation.
    4. A car exposed to the elements will rust faster in the Midwest than in Arizona.
    5. All of these are true.

 

 

 

  1. Which of the following statements is false?
    1. Stainless steel contains chromium and nickel, which form protective oxide coatings.
    2. Galvanized steel is coated with zinc to form an oxide coating.
    3. Cathodic protection is a method used to protect steel in buried tanks and pipelines.
    4. Chromium and tin are often used to plate steel by forming a durable oxide coating.
    5. All of these are true.

 

 

               

 

  1. How many faradays are involved in conversion of a mole of    to            ?
    1. 1
    2. 2
    3. 3
    4. 4
    5. 5

 

 

 

 

  1. How many moles of electrons are produced from a current of 14.4 A in 3.20 hours? A) 4.78 ´ 10–4 mol
  1. 1.72 mol
  2. 46.1 mol
  3. 3.35 mol

E)     9.33 ´ 103 mol

 

 

               

 

  1.  

4

A common car battery consists of six identical cells each of which carries out the reaction: Pb + PbO2 + 2HSO + 2H+ ® 2PbSO4 + 2H2O

Suppose that in starting a car on a cold morning a current of 125 amperes is drawn for 18.4 seconds from a cell of

the type described above. How many grams of Pb would be consumed? (The atomic weight of Pb is 207.19.) A) 4.94 g

B)     2.47 g

C)     7.29 ×10- 3 g

D)    1.58 ×10- 4 g

E)     1.19 ×10- 2 g

 

 

               

 

 

 

  1. If oxidation of H2O occurs at the anode, how many moles of oxygen gas will evolve for every 151 grams of Cr(s) deposited?

A) 4.36

B)     0.726

 

C)     17.4

D)   11.6

E)     3.87

 

 

 

 

  1. If the current is 10.0 amperes, how long will it take to deposit 105 grams of Cr(s) onto the bumper? A) 5.41 h
  1. 1.35 days
  2. 54.1 min
  3. 2.02 min
  4. 2 mo, 25 days, 14 h, and 6 s

 

 

               

 

  1. Copper is electroplated from CuSO4 solution. A constant current of 3.19 amp is applied by an external power supply. How long will it take to deposit 1.00 ´ 102 g of Cu? The atomic mass of copper is 63.546.

A) 26.4 h

  1. 13.2 min
  2. 2.03 days

D)   9.57 s

E)     3.78 h

 

 

 

 

  1. What quantity of charge is required to reduce 27.8 g of CrCl3 to chromium metal? (1 faraday = 96,485 coulombs)

A)   1.69 ×104  C

B)     5.08 ×104 C

C)     8.05 ×102 C

D)   2.68 ×102 C

E)     none of these

 

 

               

 

  1. Electrolysis of a molten salt with the formula MCl, using a current of 3.86 amp for 16.2 min, deposits 1.52 g of metal. Identify the metal. (1 faraday = 96,485 coulombs)
    1. Li
    2. Na
    3. K
    4. Rb
    5. Ca

 

 

               

 

  1. If a constant current of 5.0 amperes is passed through a cell containing Cr3+ for 1.0 hour, how many grams of Cr will plate out onto the cathode? (The atomic mass of Cr is 51.996.)
    1. 9.7 g

B)     9.0 ×10- 4 g

  1. 3.2 g
  2. 29 g

 

E)     6.2 ×10- 2 g

 

 

               

 

  1. If an electrolysis plant operates its electrolytic cells at a total current of 1.8 ´ 106 amp, how long will it take to produce one metric ton (one million grams) of Mg(s) from seawater containing Mg2+?

(1 faraday = 96,485 coulombs)

    1. 1.2 h
    2. 1.2 days
    3. 37 min

D) 0.61 h

E)     0.31 year

 

 

 

 

  1. Nickel is electroplated from a NiSO4 solution. A constant current of 5.50 amp is applied by an external power supply. How long will it take to deposit 100. g of Ni? The atomic mass of Ni is 58.69.

A)   16.6 h

B)     8.30 h

  1. 18.2 min

D)   55.4 s

E)     1.09 s

 

 

 

 

  1.  

4

A solution of MnO 2– is electrolytically reduced to Mn3+. A current of 8.07 amp is passed through the solution for

15.0 minutes. What is the number of moles of Mn3+ produced in this process? (1 faraday = 96,485 coulombs) A) 0.0753

B)     0.226

C)     4.18 ×10- 4

D) 0.0251

E)     1.67 ×10- 3

 

 

               

 

  1. How many seconds would it take to deposit 18.2 g of Ag (atomic mass = 107.87) from a solution of AgNO3 using a current of 10.00 amp?

A)   3.26 ×103  s

B)     8.14 ×102  s

C)     4.88 ×103  s

D)   1.63 ×103  s

E)     5.43 ×102  s

 

 

               

 

  1. Gold (atomic mass = 197.0) is plated from a solution of chloroauric acid, HAuCl4; it deposits on the cathode. Calculate the time it takes to deposit 0.30 gram of gold, passing a current of 0.10 amperes.

(1 faraday = 96,485 coulombs)

    1. 12 hours

 

    1. 1.2 hours
    2. 24 minutes
    3. 8.2 minutes
    4. none of these

 

 

               

 

  1. An unknown metal (M) is electrolyzed. It took 74.1 s for a current of 2.00 amp to plate 0.107 g of the metal from a solution containing M(NO3)3. Identify the metal.
    1. La
    2. Bi
    3. Ga
    4. Cu
    5. Rh

 

 

               

 

  1.  

2

Gold is produced electrochemically from an aqueous solution of Au(CN) containing an excess of CN. Gold metal and oxygen gas are produced at the electrodes. How many moles of O2 will be produced during the production of

1.00 mole of gold?

A)   0.25

B)     0.50

C)     1.00

D)   3.56

E)     4.00

 

 

 

 

  1. What mass of Ti(s) may be deposited from an aqueous TiCl2 solution if a current of 2.50 A is applied to the solution for 365 s? (e°red(Ti2+/Ti) = –1.63 V, F = 96485 C/mol)

A)   0.453 g

B)     0.906 g

C)     0.369 g

D)   0.226 g

E)     0.139 g

 

 

               

 

  1. What mass of chromium could be deposited by electrolysis of an aqueous solution of Cr2(SO4)3 for 145 minutes using a constant current of 11.0 amperes?

A)   0.187 g

B)     154.7 g

C)     0.287 g

D)   25.8 g

E)     17.2 g

 

 

               

 

  1. Which of the following are incorrectly paired?
    1. Alumina – pure aluminum oxide
    2. Downs cell – electrolyzes molten sodium chloride
    3. Mercury cell – used in preventing contamination of NaOH by NaCl
    4. Hall-Heroult process – uses cryolite in production of aluminum

 

    1. All of these are correct.

 

 

               

 

 

 

  1. Which of the following used to be more precious than gold or silver, due to difficulties refining it?
    1. copper
    2. aluminum
    3. tin
    4. zinc
    5. iron

 

 

               

 

 

 

  1. What are the products of the chlor-alkali process?
    1. sodium and sodium chloride
    2. sodium chloride and chlorine
    3. sodium and chlorine
    4. sodium hydroxide and chlorine
    5. aluminum and cryolite

 

 

               

 

 

  1.  

7                                                 3

Balance the following equation: Cr2O 2– + I ® Cr3+ + IO (acid)

 

 

 

 

  1. Balance the following equation: Zn + As2O3 ® AsH3 + Zn2+ (acid)

 

 

 

  1.  

4                                        2                   3

Balance the following equation: MnO + Br ® MnO + BrO (base)

 

 

  1.  

2                                      3

Balance the following equation: Bi(OH)3 + SnO 2– ® Bi + SnO 2– (base)

 

 

 

  1. Consider a galvanic cell with a zinc electrode immersed in 1.0 M Zn2+ and a silver electrode immersed in 1.0 M Ag+. Zn2+ + 2e ® Zn       e° = –0.76 V

Ag+ + e ® Ag                          e° = 0.80 V

Which of the electrodes is the anode?

 

 

 

 

  1.  

7

Balance the following redox equation in acidic solution. What is the coefficient of the water? CH3OH(aq) + Cr2O 2-

(aq) ® CH2O(aq) + Cr3+(aq)

    1. 1
    2. 2
    3. 4
    4. 6
    5. 7

 

 

               

 

 

  1.  

7

Balance the following redox equation in acidic solution. What is the coefficient of the H+ ion? CH3OH(aq) + Cr2O 2-(aq) ® CH2O(aq) + Cr3+(aq)
    1. 1
    2. 2
    3. 3
    4. 4
    5. 8

 

 

               

 

 

  1. Balance the following redox equation in acidic solution. What is the coefficient of the water? MnO4-(aq) + H2O2(aq) ® Mn2+(aq) + O2(g)
    1. 2
    2. 4
    3. 6
    4. 8
    5. 16

 

               

 

 

  1.  

4

Balance the following redox equation in acidic solution. What is the coefficient of the H+ ion? MnO -(aq) + H2O2(aq) ® Mn2+(aq) + O2(g)

    1. 2
    2. 4
    3. 6
    4. 8
    5. 16

 

 

               

 

 

  1.  

2

2

Balance the following redox equation in basic solution. What is the coefficient of the water? NO -(aq) + Al(s) ® NH3(g) + AlO -(aq)

    1. 1
    2. 2
    3. 4
    4. 6
    5. 7

 

 

 

 

 

  1.  

2

2

Balance the following redox equation in basic solution. What is the coefficient of the hydroxide ion? NO -(aq) + Al(s) ® NH3(g) + AlO -(aq)

    1. 1
    2. 2
    3. 4
    4. 6
    5. 7

 

 

 

 

 

  1.  

4

Balance the following redox equation in basic solution. What is the coefficient of the water? MnO -(aq) + HCOOH(aq) ® Mn2+(aq) + CO2(g)

    1. 2
    2. 4
    3. 6
    4. 8
    5. 16

 

 

 

 

  1.  

4

Balance the following redox equation in basic solution. What is the coefficient of the hydroxide ion? MnO -(aq) + HCOOH(aq) ® Mn2+(aq) + CO2(g)

    1. 2
    2. 4
    3. 6
    4. 8
    5. 16

 

 

               

 

 

  1. The reduction potential to the metal of Ba2+ is -1.57 V. Given that the cell potential for the reaction 3 Ba(s) + 2 La3+ (aq) ® 3 Ba2+(aq) + 2 La(s) is 0.52 V, the reduction potential for La3+(aq) is:

A) (0.52 – (3 x -1.57)) / 2

B)     (0.52) – (-1.57)

C)     (+1.57) + (0.52)

D) 0.52 – (+1.57)

E)     (-1.57) - (0.52)

 

 

               

 

 

  1. A voltaic cell consists of a Cu2+/Cu electrode (e°red = 0.34 V) and an Au3+/Au electrode (e°red = 1.50 V). Calculate [Au3+] if [Cu2+] = 1.20 M and ecell = 1.13 V

A)   0.001 M

B)     0.002 M

C)     0.04 M

  1. 0.2 M
  2. 5.0 M

 

 

               

 

  1. The reduction potential to the metal of Ba2+ is -1.57 V and that of Cu2+ is +0.34 V. The cell potential for the reaction Ba(s) + Cu2+(aq) ® Ba2+(aq) + Cu(s) is therefore:

A) (-1.57) + (0.34)

B)     2 x (-1.57) + 2 x (0.34)

C)     (+1.57) + (-0.34)

D)   (-1.57) + (+0.34)

E)     (+1.57) + (+0.34)

 

 

 

 

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