AP Chemistry

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Chemistry 7: Electrochemistry
A.
B.
C.
Name __________________________
Balancing Redox Equations (20.2)

assign oxidation to each atom

determine oxidized atom (oxidation # increases) and
reduced atom (oxidation # decreases)

split into oxidation and reduction half-reactions

eliminate spectator ions (ion that doesn't contain atom
that changes oxidation number—often cation)

balance each half reaction
o balance atoms except O and H
o balance O, by adding H2O
o balance H, by adding H+
o balance charge, by adding e
multiply half-reactions to equalize electrons

add half-reactions together

simplify by reducing H2O and H+ and/or coefficients

this process assumes reaction takes place in acid
(H+), if in base, add an OH- for each H+ in the final
equation (combine H+ and OH- to make water)
Standard Reduction Potentials Chart (20.4 to 20.6)
1. reduction half reactions
a. listed from “greatest” electron affinity to “least”
b. 2 H+ + 2 e-  H2: Eored = Eoox = 0 V
c. Eo measured in volts, 1 V = 1 J/C
1. "o": standard conditions (25oC, 1 atm, 1 M)
2. not proportional to amount of chemical
d. oxidation is reverse (Eoox = -Eored)
2. Eo = Eored + Eoox
a. Eo > 0 is a spontaneous reaction
(reduction listed above oxidation on chart)
b. voltage under nonstandard conditions
1. Nernst equation: E = Eo – (RTo/nF)lnQ
R (8.31), To (298) and F (96,500) are constant
(E = Eo – (0.0257/n)lnQ)
2. Q (quotient) = Product/Reactants
a. partial pressure (atm) gases,
concentration (M) of ions
b. solids and liquids excluded
Voltaic (Galvanic) Cell (20.3)
spontaneous redox reaction generates voltage  electrons
flow through wires from oxidation cell to reduction cell
anode
(–)
Voltage > 0
salt bridge
cathode
(+)
porous membrane
site of oxidation
site of reduction
1. oxidation half cell (– anode)
a. reducing agent gives up electrons to external
circuit (wires)
b. anions flow toward anode through salt bridge/
porous membrane to maintain electrical neutrality
2. reduction half cell (+ cathode)
a. oxidizing agent attract electrons from wire
b. cations flow toward cathode through salt bridge
3. predict how change affects voltage
a. reactant: [ions] or Pgases  E (voltage)
b. over time: reactant & product  E
c. size of electrode and chamber: no change
d. remove salt bridge: E = 0
D.
Electrolytic Cell (20.9)
battery forces non-spontaneous redox reaction by pulling
electrons from reducing agent and sending to oxidizing agent
anode
(+)
+ Battery –
cathode
(–)
site of oxidation
site of reduction
1. Eo < 0 (battery makes up for deficit)
2. oxidation at + anode, reduction at – cathode
3. electrolysis in water solutions (inert electrodes)
a. cathode reduction: H2O or cation (greater Eored)
1. columns 1, 2 or Al3+: 2 H2O + 2 e-  H2 + 2 OH2. acid (H+): 2 H+ + 2 e-  H2
3. otherwise: Mx+ + X e-  M
b. anode oxidation: anion or H2O
1. Cl-, Br-, I-: 2 X-  X2 + 2 e2. base (OH-): 4 OH-  O2 + 2 H2O + 4 e3. otherwise: 2 H2O  O2 + 4 H+ + 4 e4. electroplating (transition metal cations coat cathode)
a. current, I, measured in amperes (amps—A)
1 A = 1 C/s (coulomb/second)
b. mass plated given current, I, and time, t
(t) s x (I) C x mol e- x mol Mx+ x (MM) g = __ g
1 s 96,500 C X mol e- 1 mol Mx+
c. time for plating—calculate right to left
(2) Given that pink indicates the presence of OH-,
write the half-reaction that occurs at the cathode.
Experiments
1.
Voltaic Cell Lab—Measure the voltage of six voltaic cells,
determine the reduction half-cell potential for each metal
and compare these results to the standard values.
Polish the metal electrodes (1-Ag, 2-Cu, 3-Pb, and 4-Zn).
Half fill each quadrant of the Petri dish with the 0.1 M nitrate
salt of the corresponding metal. Drape the KNO3-soaked
ring so that it is submerged in each quadrant. Measure the
voltage of each system (reverse the electrodes when V < 0).
Record the metal that is connected to the black electrode
(oxidation) and the voltage (Elab) for the overall reaction.
a. Complete the chart for each voltaic cell.
Ag and Cu
Ag and Pb
Ag and Zn
Elab
Elab
Elab
Black
Black
Black
(1) What color change occurs at the anode?
(2) Given that amber indicates the presence of I2,
write the half-reaction that occurs at the anode.
c.
Are these reactions consistent with the predicted
reactions based on KI as the electrolyte? Explain
Ag + Cu
Part 2: Measure the change in mass of the zinc electrode
and the volume of hydrogen gas produced during
electrolysis, calculate the molar mass of zinc and compare
the value to the periodic table.
Fill a 150 mL beaker ¾ full with conducting solution. Fill the
25 mL volumetric flask with the conducting solution from the
beaker. Place rubber dam over the mouth of the flask, invert
it and place it in the beaker (mouth down). Scrap off the
rubber dam and slip the J-shaped electrode inside the flask
without getting any air in. Polish the zinc and mass it (m1).
Connect the wire lead attached to the black side of the
power supply to the J-electrode (cathode). Connect the wire
lead attached to the red side of the power supply to the zinc
(anode). Turn on the power supply. Hydrogen gas should
be bubbling from the J-electrode. Record the start time (t1).
When the water level in the flask is near the 25 mL line,
carefully raise the flask so that the level of the conducing
fluid solution inside the flask is the same as the level in the
beaker. When the water line in the flask is at 25 mL, turn the
power supply off and disconnect the wires from the zinc and
J-electrode. Record the stop time (t2). Return the conducting
solution to the stock bottle. Scrub the zinc with a scrub pad
under running water. Thoroughly dry the zinc and then
mass it (m2). Record the temperature (T), lab pressure (Plab)
and water vapor pressure (PH2O).
d. Record the data.
m1
t1
m2
t2
V
T
Plab
PH2O
25.0 mL
e. Complete and balance the redox reaction between zinc
metal and water.
Reduction
Ag + Pb
Oxidation
Ag + Zn
d. Compare the voltages of the remaining systems with
the table values. (There is no need to correct for 0.1 M)
Eolab
Eotable = Eored + Eoox
Cell
%
Overall
f. Calculate the molar mass of Zn by completing the chart.
Cu + Pb
Volume in L
Cu and Pb
Elab
Black
b.
Ag
+
Cu
Cu and Zn
Elab
Black
Pb and Zn
Elab
Black
Determine the standard voltage (Eolab) for each cell.
Half Reaction
Reduction
Oxidation
Overall
Eolab
Ag
+
Pb
Reduction
Oxidation
Overall
Eolab
Ag
+
Zn
Reduction
Oxidation
Overall
Eolab
c.
Determine the lab values for Eored for Cu, Pb and Zn
given Eored for Ag is 0.80.
Eored = 0.80 – Eolab
Eotable
Cell
%
Cu + Zn
2.
b.
Pb + Zn
Electrolysis Lab—Part 1: Observe the electrolysis of KI(aq).
Fill a weighing boat with 0.5 M KI. Add 3 drops indicator.
With the power supply OFF, attach a piece of graphite to
each wire lead and then plug in the power supply into the
electric outlet. Dip the two pieces of graphite into opposite
sides of the weighing boat for 10 seconds. Observe
changes at the cathode (black electrode) and anode (red
electrode), record the observations. Turn off the power
supply and remove the graphite.
a. (1) What color change occurs at the cathode?
Mass of Zinc lost
Pressure in atm
Temperature in K
Moles of H2
Molar Mass of Zinc
%
g.
Why is it necessary to equalize the water level inside
the volumetric flask with the water level in the beaker?
h.
Calculate the average current during the electrolysis
by completing the chart.
8.
Total charge Q in C
Total time T in s
The mass percent of H2O2 in a hydrogen peroxide solution is
determined by titration with an acidified solution of KMnO4.
a. Balance the redox reaction:
MnO4- + H2O2 + H+  Mn2+ + O2 + H2O.
oxidation
reduction
Current I (I = Q/t)
i.
overall
b. How many moles of MnO4- are needed to react all the
H2O2 if 2.647 mL of 0.0200 M KMnO4 is added?
Calculate the approximate voltage needed to perform
the electrolysis. (Assume standard conditions)
Practice Problems
1.
2.
A. Oxidation-Reduction Reactions
Balance the redox equation in acid.
MnO4- + NO2- + H+  Mn2+ + NO3- + H2O
Balanced the redox equation in base
Cr(OH)3(s) + ClO-  Cl- + CrO42-
3.
Balance the redox equation in acid:
Zn(s) + NO3-  Zn2+ + NH4+.
4.
Balanced the redox equation in base:
NO2- + Al(s)  NH3(aq) + Al(OH)4-
5.
Balance the disproportionation reaction (same element
undergoes oxidation and reduction): Cl2(g)  Cl- + ClO-.
in
acid
How many moles of H2O2 reacted with the MnO4-?
d.
How many grams of H2O2 reacted with the MnO4-?
e.
What is the % H2O2 if 0.150 g of solution is used?
B. Standard Reduction Potentials Chart
Consider the Standard Reduction Potential Chart.
a. What are the standard conditions?
Highlight the correct option.
b. When comparing two reactants (on the left side of the
chart), the reactant that is listed (higher/lower) on the
chart is the stronger oxidizing agent.
c. When comparing two products (on the right side of the
chart), the product that is listed (higher/lower) on the
chart is the stronger reducing agent.
d. The strongest reducing agents are found in column
(1/17) on the periodic table.
e. The strongest oxidizing agents are found in column
(1/17) on the periodic table.
f. Reactive (metals/nonmetals) tend to lose electrons
and act as good (oxidizing/reducing) agents.
g. When combining two half reactions, the left side agent
listed higher on the chart will (give/receive) electrons
from the right side agent listed lower on the chart.
h. Write the equations for each half-reaction and overall
reaction. Calculate Eo.
Eo
Balanced reaction
Zn  Zn2+
in
base
6. Balance the reaction: P4(s)  H2PO2- + PH3(g)
Br2  BrOverall
H2  H+
in
acid
Ag+  Ag
in
base
7. Write net ionic equations for the following redox reactions.
a. Solid iron + iron(III) sulfate
b.
9.
c.
Potassium dichromate + acidified hydrogen peroxide.
Overall
i.
j.
A spontaneous reaction has a (positive/negative) Eotot.
Is the redox reaction, Cu(s) + Cl2(g)  CuCl2(s),
spontaneous? Support your answer with calculations.
10. Indicate where on the Standard Reduction Potentials Chart
you would find:
a. The chemical species that is the easiest to oxidize.
b.
The chemical species that is the easiest to reduce.
11. Using standard reduction potentials, calculate the standard
voltage for each of the following reactions.
Cl2(g) + 2 I -  2 Cl- + I 2(s)
Ni(s) + 2 Fe3+  Ni2+ + Fe2+
Fe(s) + 2 Fe3+  2 Fe2+
2 Al3+ + 3 Ca(s)  2 Al(s) + 3 Ca2+
12. Using standard reduction potentials, highlight the stronger
reducing agents for each of the following pairs.
Fe(s) or Mg(s) Ca(s) or Al(s) H2(g) or H2S(g) Sn2+ or Fe2+
13. The unbalanced reduction half-reactions that operate in a
car battery follow.
PbSO4(s)  Pb(s) + SO42Eored = -0.356 V
2PbO2(s) + SO4  PbSO4(s)
Eored = 1.687 V
a. Write balanced half reactions and the overall reaction.
Label which reaction occurs at the anode and cathode.
b.
Calculate Eotot for the cell.
c.
Write the Nernst equation for the overall reaction.
Calculate the voltage when [H+] = 2.0 x 10-4 M and
[SO42-] = 0.50 M.

e.
e- 

C. Voltaic (Galvanic) Cell
14. Answer the following questions based on the diagram.
e- 

If you made a voltaic cell out of this, what half-reaction
would be occurring at the cathode, and what halfreaction would be occurring at the anode?

c.
What would happen to the voltage of the battery if the
concentration of [H+] increased justify your answer)?
Voltage =
oxidation half reaction
reduction half reaction
overall reaction
16. For the generic reaction: A + B  A- + B+, for which Eo is a
positive number, answer the following questions:
a. What is being oxidized, and what is being reduced?
b.

d.
15. Complete the voltaic cell drawing using the half-reactions.
Ni2+ + 2 e-  Ni(s)
Eored = -0.25 V
Fe2+ + 2 e-  Fe(s)
Eored = -0.44 V
Label the anode metal, cathode metal, site of oxidation,
site of reduction, cation flow, anion flow, voltage, oxidation
half-reaction, reduction half-reaction, and overall reaction.
(You will need to figure out which reaction occurs at the
anode, which at the cathode, and what the voltmeter would
read under standard conditions. Assume each metal
electrode is immersed in 1 M nitrate salt of the metal.)
Which half-reaction from (b) is higher on the Standard
Reduction Potential Chart?

17. A voltaic cell consists of a strip of aluminum in a solution of
Al(NO3)3 in one beaker, and in the other beaker a strip of
nickel in a solution of Ni(NO3)2. The overall reaction is:
2 Al(s) + 3 Ni2+  2 Al3+ + 3 Ni(s)
a. What is being oxidized, and what is being reduced?

b.
Write the half-reactions that occur in the beakers.
(Indicate which reaction takes place at the anode and
which takes place at the cathode.)

c.
Indicate the signs of the two electrodes.

a. which
Zn Zn2+ Cu Cu2+
species is reduced?
species is oxidized?
species gives up its electron?
species accepts electrons?
ion passes through salt bridge?
reducing agent?
oxidizing agent?
electrode is the cathode?
electrode is the anode?
oxidation half cell
reduction half cell
b. What memory device can be used to remember that
the cathode is the site of reduction and + ions pass
through the porous membrane to the cathode?
d.
Do electrons flow from the aluminum to the nickel or
from the nickel to the aluminum?
e.
In which directions do the cations migrate and in which
direction do the anions migrate through the solution?
18. A voltaic cell consists of a strip of lead metal in a solution
of Pb(NO3)2 in one beaker, and in the other beaker a
platinum electrode is immersed in a NaCl solution, with Cl 2
gas bubbled around the electrode. The two beakers are
connected with a salt bridge.
a. Write the equation for the overall cell reaction.
b.
What is the overall voltage generated by the cell under
standard conditions?
c.
Which electrode serves as the anode, and which
serves as the cathode?
22. An electrolytic cell contains a solution of Cr(NO3)3.
a. Write the anode, cathode and overall equations.
anode
d.
Does the Pb electrode gain or lose mass as the cell
reaction proceeds?
cathode
overall
b. How long will it take to deposit 15.0 g of chromium
metal, using a current of 4.50 A?
D. Electrolytic Cell
19. Answer the following questions based on the electrolysis of
fused (melted) sodium chloride.
c.
A current of 4.50 A for 30.0 minutes passed through
the cell. The initial electrolyte contained 250 mL of
1.00 M Cr(NO3)3. Determine the
(1) Initial moles of Cr3+.
(2) moles of Cr3+ reacted.
(3) Concentration of Cr3+ after 30.0 min. of electrolysis.
a. which
Cl2 Cl- Na
species is reduced?
species is oxidized?
species gives up its electron?
species accepts electrons?
species has a higher electron affinity?
species is the reducing agent?
species is the oxidizing agent?
b. Write the equations for the electrolysis of NaCl.
oxidation half reaction
(4) Concentration of H+ after 30.0 min. of electrolysis.
Na+
reduction half reaction
overall reaction
c. What mass of Na(l) is produced using a current of
3.00 A for one hour?
20. Which cell...
Electrolytic
Voltaic
a. has a battery?
b. has a salt bridge?
c. is spontaneous?
d. has a positive anode?
21. Write equations for the oxidation, reduction and overall
reactions for the electrolysis of the salt solutions.
(5) Mass of Cr(s) that plate out at the cathode.
Practice Multiple Choice
Briefly explain why the answer is correct in the space provided.
1.
H2Se + 4 O2F2  SeF6 + 2 HF + 4 O2
Which is true regarding the reaction represented above?
(A) Oxidation number of O does not change.
(B) Oxidation number of H changes from -1 to +1.
(C) Oxidation number of F changes from +1 to -1.
(D) Oxidation number of Se changes from -2 to +6.
2.
6 I- + 2 MnO4- + 4 H2O  3 I2 + 2 MnO2 + OHWhich statement regarding the reaction is correct?
(A) Iodide ion is oxidized by hydroxide ion.
(B) MnO4- is oxidized by iodide ion.
(C) Manganese oxidation number changes from +7 to +2.
(D) Oxidation number of iodine changes from -1 to 0.
3.
2 H2O + 4 MnO4- + 3 CIO2-  4 MnO2 + 3 CIO4- + 4 OHWhich species acts as an oxidizing agent in the reaction?
(A) MnO4- (B) CIO4(C) CIO2(D) MnO2
4.
Which species CANNOT function as an oxidizing agent?
(A) Cr2O72- (B) MnO4- (C) NO3(D) I-
5.
When acidified solutions of K2Cr2O7 and Na2S are mixed,
Cr3+ and S are formed. Which is the best reducing agent?
(A) K2Cr2O7 (B) Na2S
(C) Cr3+
(D) S
6.
_Ag+ + _AsH3 + _OH-  _Ag + _H3AsO3 + _H2O
What is the coefficient for OH- in the balanced equation?
(A) 2
(B) 4
(C) 5
(D) 6
NaCl
CuSO4
Ba(OH)2
HNO3
Na2CO3
KF
7.
_Cr2O72- + _e- + _H+  _Cr3+ + _H2O
What is the coefficient for H+ in the balanced half-reaction?
(A) 2
(B) 6
(C) 7
(D) 14
18. 10 HI + 2 KMnO4 + 3 H2SO4  5 I2 + 2 MnSO4 + K2SO4 + 8 H2O
How many moles of HI are used to produce 2.5 mol of I2?
(A) 5.0
(B) 8.0
(C) 10.
(D) 12
8.
_Mg + _NO3– + _H+  _Mg2+ + _NH4+ + _H2O
What is the coefficient for H+ in the balanced equation?
(A) 1
(B) 3
(C) 5
(D) 10
19. 2 H2O + 4 MnO4- + 3 CIO2-  4 MnO2 + 3 CIO4- + 4 OHHow many moles of ClO2- react with 0.20 L of 0.20 M MnO4-?
(A) 0.030 (B) 0.053 (C) 0.075 (D) 0.13
9.
_CrO2– + _OH–  _CrO42– + _H2O + _e–
What is the ratio of the coefficients OH–/CrO2– in the
balanced half-reaction
(A) 1:1
(B) 2:1
(C) 3:1
(D) 4:1
20.
10.
_Fe(OH)2 + _O2 + _H2O  _Fe(OH)3
If 1 mole of O2 oxidizes Fe(OH)2, how many moles of
Fe(OH)3 can be formed?
(A) 2
(B) 3
(C) 4
(D) 5
11. In which species does sulfur have the same oxidation
number as it does in H2SO4?
(A) H2SO3 (B) S2O32- (C) S2(D) SO2Cl2
12.
2 HClO + 3 O2  2 HClO4
As the reaction proceeds to the right, the oxidation number
of chlorine changes from
(A) -1 to +3 (B) -1 to +5 (C) +1 to +7 (D) +3 to +7
13. Which will generate H2(g) when added to 1 M HCl?
(A) CuS
(B) Zn
(C) CaCO3 (D) Mg(OH)
14.
_Cr2O72- + _H2S + _H+  _Cr3+ + _S + _H2O
What is the coefficient for H+ in the balanced equation?
(A) 2
(B) 4
(C) 6
(D) 8
15.
3 Cu + 8 H+ + 2 NO3-  3 Cu2+ + 2 NO + 4 H2O
Which statements about the reaction are true?
I. Cu acts as an oxidizing agent.
II. Nitrogen's oxidation state changes from +5 to +2.
III. Hydrogen ions are oxidized to form H2O.
(A) I only
(B) II only (C) III only (D) I and II
16. In which reaction does the same element undergo both
oxidation and reduction?
(A) S8(s) + 8 O2(g)  8 SO2(g)
(B) 3 Br2(aq) + 6 OH-  5 Br- + BrO3- + 3 H2O
(C) Ca2+ + SO42-  CaSO4(s)
(D) PtCI4(s) + 2 CI-  PtCI62-
17. Which reaction is an oxidation-reduction reaction?
(A) HC2H3O2(aq) + NH3(aq)  C2H3O2- + NH4+
(B) Ba2+ + SO42-  BaSO4(s)
(C) Zn(OH)2(s) + 2 OH-  [Zn(OH)4]2(D) 2 K(s) + Br2(l)  2 KBr(s)
5 Fe2+ + MnO4- + 8 H+  5 Fe3+ + Mn2+ + 4 H2O
25.0 mL of an acidified Fe2+ solution requires 14.0 mL of
0.10-M MnO4- solution to reach the equivalence point. The
concentration of Fe2+ in the original solution is
(A) 0.10 M (B) 0.56 M (C) 0.28 M (D) 0.14 M
21. Use the reduction potentials to determine which one of the
reactions below is spontaneous.
-0.5 V
0.5 V
Cd2+ + 2 e-  Cd
Cu+ + 1 e-  Cu
2+
Mn + 2 e  Mn -1.2 V
Fe3+ + 1 e-  Fe2+ 0.7 V
(A) Cd2+ + 2 Cu  Cd + 2 Cu
(B) Mn2+ + 2 Cu  Mn + 2 Cu+
(C) Cd2+ + Mn  Cd + Mn2+
(D) Cu+ + Fe3+  Cu + Fe2+
22. According to the information below, what is the standard
reduction potential for the half-reaction: M3+ + 3 e-  M?
M + 3 Ag+  3 Ag + M3+
Eo = + 2.5 V
+
Ag + e  Ag
Eo = + 0.8 V
(A) -1.7 V (B) -0.1 V (C) 0.1 V
(D) 1.7 V
23. Magnesium reacts with dilute hydrochloric acid to produce
hydrogen gas. Silver does not react in dilute hydrochloric
acid. Based on this information, which of the following
reactions will occur spontaneously?
(A) H2(g) + Mg2+  2 H+ + Mg(s)
(B) 2 Ag(s) + Mg2+  2 Ag+ + Mg(s)
(C) 2 Ag+ + Mg(s)  2 Ag(s) + Mg2+
(D) 2 Ag + 2 H+  H2(g) + 2 Ag+
24.
Zn + Cu2+  Zn2+ + Cu
Which could account for the observed voltage of 1.00 V
instead of the standard cell potential, Eo, of 1.10 V?
(A) The copper electrode was larger than the zinc electrode.
(B) The Zn2+ electrolyte was Zn(NO3)2, while the Cu2+
electrolyte was CuSO4.
(C) The concentration of the Zn2+ solution was greater
than the Cu2+ solution.
(D) The solutions in the half-cells had different volumes.
Questions 25-26 refer to an electrolytic cell that involves the
following half-reaction. AIF63- + 3 e-  Al + 6 F25. Which of the following occurs in the reaction?
(A) AIF63- is reduced at the cathode.
(B) Al is oxidized at the anode.
(C) Aluminum is converted from the -3 oxidation state to
the 0 oxidation state.
(D) F- acts as a reducing agent.
26. A steady current of 10 A is passed through an aluminumproduction cell for 15 minutes. Which of the following is
the correct expression for calculating the number of grams
of aluminum produced? (1 faraday = 96,500 coulombs)
(A) (10)(15)(96,500)/(27)(60)
(B) (10)(15)(27)/(60)(96,500)
(C) (10)(15)(60)(27)/(96,500)(3)
(D) (96,500)(27)/(10)(15)(60)(3)
27. Which of the following expressions is correct for the
maximum mass of copper, in grams, that could be plated
out by electrolyzing aqueous CuCl2 for 16 hours at a
constant current of 3.0 A? (1 faraday = 96,500 coulombs)
(A) (16)(3,600)(3.0)(63.55)(2)/(96,500)
(B) (16)(3,600)(3.0)(63.55)/(96,500)(2)
(C) (16)(3,600)(3.0)(63.55)/(96,500)
(D) (16)(60)(3.0)(96,500)(2)/(63.55)
34. A power supply has lost the markings that indicate the
positive and negative. A chemist suggests that the
terminals be connected to a pair of platinum electrodes
that dip into 0.1 M KI solution. Which correctly identifies
the polarities of the terminals?
(A) A gas will be evolved only at the positive electrode.
(B) A gas will be evolved only at the negative electrode.
(C) An amber color will appear near the negative electrode.
(D) A metal will be deposited on the positive electrode.
35. In the electroplating of nickel, 0.200 faraday of electrical
charge is passed through a solution of NiSO4. What mass
of nickel is deposited?
(A) 2.94 g (B) 5.87 g (C) 11.7 g (D) 58.7 g
Practice Free Response
28. If 0.060 faraday is passed through an electrolytic cell
containing a solution of In3+ ions, the maximum number of
moles of In that could be deposited at the cathode is
(A) 0.010 (B) 0.020 (C) 0.030 (D) 0.060
1.
Questions 29-33 The spontaneous reaction that occurs when
the cell below operates is 2 Ag+ + Cd  2 Ag + Cd2+
A power supply is connected to two platinum electrodes
immersed in a beaker containing 1.0 M CuSO4(aq) at 25oC.
As the cell operates, copper metal is deposited on the left
electrode and O2(g) is produced at the right electrode. The
reduction half-reactions that occur are:
O2(g) + 4 H+(aq) + 4 e-  2 H2O(l) Eo = +1.23 V
Cu2+(aq) + 2 e-  Cu(s)
Eo = +0.34 V
a. Is the direction of electron flow in the wire from left to
right or from right to left?
b.
Write a balanced net ionic equation for the electrolysis
reaction that occurs.
A current of 1.50 A passes through the cell for 40.0 minutes.
c. Calculate the mass of the Cu(s) that is deposited.
Which occurs for each of the following circumstances?
(A) Voltage increases.
(B) Voltage decreases but remains above zero.
(C) Voltage becomes zero and remains at zero
(D) No change in voltage occurs
29. A 50-mL sample of a 2-M Cd(NO3)2 solution is added to
the left beaker.
30. The silver electrode is made larger.
31. The salt bridge is replaced by a platinum wire.
d.
2.
Calculate the dry volume, in liters, measured at 25oC
and 1.16 atm, of the O2(g) that is produced.
A voltaic cell is constructed with a strip of Sn in a solution of
Sn(NO3)2 in one container and a strip of an unknown metal,
X, in a solution of X(NO3)3 in another container. The two
containers are connected by a salt bridge and the two metal
strips are connected with a metal wire. The mass of the Sn
electrode increases. The half-reactions are:
Sn2+(aq) + 2 e-  Sn(s)
Eo = -0.14 V
3+
X (aq) + 3 e  X(s)
Eo = ?
a. Which electrode is the cathode? Justify your answer.
32. Current is allowed to flow for 5 minutes.
b.
What directions do electrons flow in terms of the Sn
and unknown metal strips.
33. The silver electrode is replaced by a copper electrode.
c.
If the cell potential, Eocell, is + 0.60 V, what is the
standard reduction potential for X3+(aq) + 3 e-  X(s)?
d.
Identify metal X from the chart of standard potentials.
3.
e.
Write a balanced net-ionic equation for the overall
chemical reaction occurring in the cell.
f.
If the [Sn2+] = 0.50 M and [X3+] = 0.10 M. What is the
cell potential, Ecell?
4.
An electrochemical cell is constructed with an open switch,
as shown in the diagram above. A strip of Sn and a strip of
an unknown metal, X, are used as electrodes. When the
switch is closed, the mass of the Sn electrode increases.
An electrolytic cell contains a solution of Cr(NO3)3. Assume
that chromium metal plates out at one electrode and
oxygen gas is evolved at the other electrode.
a. Write the anode half-reaction, the cathode halfreaction, and the overall reaction for the cell.
b.
How long will it take to deposit 15.0 g of chromium
metal, using a current of 4.50 A?
c.
A current of 4.50 A is passed through the cell for 30.0
min. If we start out with 250 mL of 1.00 M Cr(NO3)3,
what is the concentration of Cr3+ after electrolysis?
A voltaic cell consists of two half-cells, one of which
contains a Pt electrode surrounded by Cr3+ and Cr2O72ions. The other half-cell contains a Pt electrode surrounded
by Mn2+ ions and MnO2(s). Assume the cell reactions,
which produce a positive voltage, involves both Cr3+ and
Mn2+ ions.
Cr2O72- + 14 H+ + 6 e-  2 Cr3+ + 7 H2O
Eored = 1.33 V
+
2+
MnO2(s) + 4 H + 2 e  Mn + 2 H2O
Eored = 1.229 V
d. Write the anode half-reaction, the cathode halfreaction, and the overall reaction for the cell.
e.
Calculate Eotot for the cell.
f.
Calculate the voltage of the cell when all ionic species
except H+ have a concentration of 0.300 M and the
solution has [H+] = 3.16 x 10-4 M.
The half-reactions are shown below.
Sn2+ + 2 e-  Sn(s)
Eo = -0.14 V
X3+ + 3 e-  X(s)
Eo = ?
a. In the diagram above, label the electrode that is the
cathode. Justify your answer.
b.
In the diagram above, draw an arrow indicating the
direction of the electron flow in the external circuit
when the switch is closed.
c.
If the standard cell potential, Eocell, is + 0.60 V, what is
the standard reduction potential, in volts, for the X3+/X
electrode?
d.
Identify metal X.
e.
Write a balanced net-ionic equation for the overall
chemical reaction occurring in the cell.
f.
In the cell, the concentration of Sn2+ is changed from
1.0 M to 0.50 M, and the concentration of X3+ is
changed from 1.0 M to 0.10 M.
(1) Substitute all the appropriate values for
determining the cell potential, Ecell, into the Nernst
equation. (Do not do any calculations.)
(2) On the basis of your response in part (f) (1), will
the cell potential, Ecell, be greater than, less than,
or equal to the original Eocell? Justify your answer.
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