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AP Chemistry 9: Electrochemistry
A.
Oxidation-Reduction Reactions (4.4, 20.1 to 20.2)
1. reactions involve the transfer of electrons (or control of
electrons) from substance that is oxidized (reducing
agent) to substance that is reduced (oxidizing agent)
Electrons
Process
Agent
Lose
Oxidation
Reducing Agent
Gain
Reduction
Oxidizing Agent
2. oxidation number
a. each atom is assigned an oxidation number which
represents the number of electrons (compared to
a neutral atom) that the atom has control over
1. zero for isolated neutral atom
2. equals ionic charge for monatomic ions
b. overall oxidation number for polyatomic species
1. zero for neutral compound or molecule
2. equals ionic charge for polyatomic ion
c. assign oxidation numbers to atoms within a
molecule or polyatomic ion
1. assign the standard value for the following
a. Li+,...(+1), Mg2+...(+2), Al (+3), F- (-1)
b. O (-2) except for peroxides, O22- (-1)
c. H (+1) except for hydrides, MHx (-1)
2. determine the missing atom's value by using
the total for the compound (see 2b)
3. balancing redox equations

assign oxidation to each atom (see 2c)

determine oxidized atom (oxidation # increases) and
reduced atom (oxidation # decreases)

split the reaction into an oxidation half-reaction and a
reduction half-reaction

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
multiple 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)
4. reduction half reactions of common oxidizing agents
a. MnO4- + 8 H+ + 5 e-  Mn2+ + 4 H2O
b. Cr2O72- + 14 H+ + 6 e-  2 Cr3+ + 7 H2O
5. oxidation of reactive metals
a. metal ion takes electron from metal listed lower on
the Standard Potential Chart (see section B)
metal + salt  salt of metal + new metal
Zn(s) + Cu(NO3)2(aq)  Zn(NO3)2(aq) + Cu(s)
net ionic: Zn(s) + Cu2+  Zn2+ + Cu(s)
b. H+ ion takes electron from metal with negative
standard potential (see chart in section B)
metal + acid  salt of metal + hydrogen gas
Zn(s) + 2 HCl(aq)  ZnCl2(aq) + H2(g)
net ionic: Zn(s) + 2 H+  Zn2+ + H2(g)
c. hydrogen in water takes electron from column 1
and heavier column 2 metals
2 Na(s) + 2 H2O  2 Na+ + 2 OH- + H2(g)
Ca(s) + 2 H2O  Ca2+ + 2 OH- + H2(g)
6. redox with metal hydride and water (the hydrogen in the
hydride (ox # = -1) reacts with the hydrogen in water (ox
# = +1) to form H2 (ox # = 0)
NaH(s) + H2O  Na+ + OH- + H2(g)
CaH2(s) + 2 H2O  Ca2+ + 2 OH- + 2 H2(g)
Name __________________________
B.
Standard Reduction Potentials Chart (20.4 to 20.6)
STANDARD REDUCTION POTENTIALS (Eo)
IN AQUEOUS SOLUTIONS
25oC, 1 M for ions, 1 atm for gases
Eo
–
–
F2(g) + 2 e
 2F
2.87
 Co2+
Co3+ + 1 e–
1.82
 Au(s)
Au3+ + 3 e–
1.50
Cl2(g) + 2 e–
 2 Cl–
1.36
O2(g) + 4 H+ + 4 e–  2 H2O(l)
1.23
Br2(l) + 2 e–
 2 Br–
1.07
 Hg22+
2 Hg2+ + 2 e0.92
 Hg(l)
Hg2+ + 2 e–
0.85
 Ag(s)
Ag+ + 1 e–
0.80
 2 Hg(l)
Hg22+ + 2 e–
0.79
 Fe2+
Fe3+ + 1 e–
0.77
 2 I–
0.53
I2(s) + 2 e–
 Cu(s)
Cu+ + 1 e–
0.52
 Cu(s)
Cu2+ + 2 e–
0.34
 Cu+
Cu2+ + 1 e–
0.15
 Sn2+
Sn4+ + 2 e–
0.15
 H2S(g)
S(s) + 2 H+ + 2 e–
0.14
 H2(g)
2 H+ + 2 e–
0.00
 Pb(s)
Pb2+ + 2 e–
–0.13
 Sn(s)
Sn2+ + 2 e–
–0.14
 Ni(s)
Ni2+ + 2 e–
–0.25
 Co(s)
Co2+ + 2 e–
–0.28
 Tl(s)
Tl+ + 1 e–
–0.34
 Cd(s)
Cd2+ + 2 e–
–0.40
 Cr2+
Cr3+ + 3 e–
–0.41
 Fe(s)
Fe2+ + 2 e–
–0.44
 Cr(s)
Cr3+ + 3 e–
–0.74
 Zn(s)
Zn2+ + 2 e–
–0.76
 H2(g) + 2 OH– –0.83
2 H2O(l) + 2 e–
 Mn(s)
Mn2+ + 2 e–
–1.18
 Al(s)
Al3+ + 3e–
–1.66
 Be(s)
Be2+ + 2 e–
–1.70
 Mg(s)
Mg2+ + 2 e–
–2.37
 Na(s)
Na+ + 1 e–
–2.71
 Ca(s)
Ca2+ + 2 e–
–2.87
 Sr(s)
Sr2+ + 2 e–
–2.89
 Ba(s)
Ba2+ + 2 e–
–2.90
 Rb(s)
Rb+ + 1 e–
–2.92
 K(s)
K+ + 1 e–
–2.92
 Cs(s)
Cs+ + 1 e–
–2.92
 Li(s)
Li+ + 1 e–
–3.05
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. Go = –nFEo (in joules)
1. n: # e- in balanced redox equation
2. F: faraday = 96,500 C/mol ec. voltage under nonstandard conditions
1. Nernst equation: E = Eo – (RTo/nF)lnQ
R (8.31), To (298) and F (96,500) are constant
2. Q (quotient) = Product/Reactants
a. partial pressure (atm) gases,
concentration (M) of ions
b. solids and liquids excluded
C.
Voltaic (Galvanic) Cell (20.3)
spontaneous redox reaction generates voltage  electrons
flow through wires from oxidation cell to reduction cell
anode
(–)
cathode
(+)
Voltage > 0
salt bridge
porous membrane
site of oxidation
site of reduction
1.
D.
oxidation half cell (– anode)
a. reducing agent (|lower| electron affinity) 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 (|higher| electron affinity) attract
electrons from external circuit (wires)
b. cations flow toward cathode through salt bridge/
porous membrane to maintain electrical neutrality
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
Electrolytic Cell (20.9)
battery forces non-spontaneous redox reaction by pulling
electrons from reducing agent and sending to oxidizing agent
anode
(+)
site of oxidation
1.
2.
3.
4.
+ Battery –
cathode
(–)
site of reduction
Eo < 0 (battery makes up for deficit)
oxidation at + anode, reduction at – cathode
electrolysis in water solutions (inert electrodes)
a. cathode reduction: H2O or cation (which ever one is
higher on the standard potential chart)
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 eelectroplating (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
b. Determine the lab values for Ered for Cu, Pb and Zn
given Ered for Ag by completing the chart.
(1) Write the equation for the oxidation half-reaction
(black) and overall reaction for each cell.
(2) Fill in the voltages (from a) as the overall Elab.
(3) Calculate Eox for Cu, Pb and Zn.
Elab
Half Reaction
Oxidation
Experiments
1.
Percent Peroxide Lab—Measure the volume of KMnO4(aq)
used to oxidize (H2O2) in order to determine the percent
peroxide in solution.
Mass the 0.0200 M KMnO4 bottle and the 3 % H2O2 bottle.
Add 3 drops of 3% H2O2, 5 drops of 3 M H2SO4 and 10 mL
distilled water to a clean 50 mL beaker. Add the KMnO4
drop by drop until a pink color persists. Mass the KMnO4
bottle and H2O2 bottle. Repeat two more times without
emptying the beaker.
a, Record the data and calculate the changes in mass.
0.0200 M KMnO4
H2O2 Solution
Trials
Start
End Change Start
End Change
1
Cu
.74
2 Ag+ + 2 e-  2 Ag(s)
.74
2 Ag+ + 2 e-  2 Ag(s)
.74
Oxidation
Pb
Reduction
Overall
Oxidation
3
b. Balance the redox reaction:
MnO4- + H2O2 + H+  Mn2+ + O2 + H2O.
oxidation
Zn
c.
reduction
d.
L MnO4-
Overall
Calculate the correction for non-standard 0.1 M
concentration for the general reaction: M  M2+ + 2 e-
Correct for the non-standard conditions and determine
the percent difference from the table.
Eotable
Elab +
E = Eolab
%
Cu
Pb
mol MnO4-
Zn
e. Compare the voltages of the remaining systems with
the table values. (There is no need to correct for 0.1 M)
Elab
Eo % 
Half Reaction
Oxidation
Cu
+
Reduction
Pb
Overall
mol H2O2
g H2O2
% H2O2
Determine the % deviation between the three trials by
completing the chart.
Trial 1
Trial 2
Trial 3
Cu
+
Zn
% H2O2
mean
Pb
+
Zn
absolute 
average 
3.
%
Voltaic Cell Lab—Measure the voltage of four voltaic cells,
determine the reduction half-cell potential for each metal
and compare these results to the standard values.
Half fill each quadrant of the Petri dish with the following: I
(0.1 M AgNO3), II (0.1 M Cu(NO3)2), III (0.1 M Pb(NO3)2)
and IV (0.1 M Zn(NO3)2). Drape KNO3-soaked strings
between quadrants (6 strings). 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 (E) 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
Cu and Pb
Elab
Black
Reduction
E = (RTo/nF)ln[M2+]
overall
c. Determine percent H2O2 by completing the chart.
Trial 1
Trial 2
Trial 3
2.
2 Ag+ + 2 e-  2 Ag(s)
Overall
2
d.
Reduction
Cu and Zn
Elab
Black
Pb and Zn
Elab
Black
Oxidation
Reduction
Overall
Oxidation
Reduction
Overall
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?
(2) Given that pink indicates the presence of OH-,
write the half-reaction that occurs at the cathode.
b.
(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?
Practice Problems
1.
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
50 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. Record the start time
(t1). Hydrogen gas should be bubbling from the J-electrode.
When the water level in the flask is near the 50 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 50 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 look up the water vapor pressure (PH2O).
d. Record the data.
m1
t1
m2
t2
V
T
Plab
PH2O
50.0 mL
e. Complete and balance the redox reaction.
Zn(s) + ___H+ 
f. Calculate the molar mass of Zn by completing the chart.
2.
A. Oxidation-Reduction Reactions
Highlight the term which correctly completes the sentences.
a. Oxidizing agents cause (oxidation/reduction) in another
atom by the process of (oxidation/reduction), which
results in the agent (gaining/losing) electrons.
b. Reducing agents cause (oxidation/reduction) in another
atom by the process of (oxidation/reduction), which
results in the agent (gaining/losing) electrons.
Determine the oxidation number of the underlined atom?
SO32MnO4H2O2
3.
Cr2O72Balance the redox equation in acid.
MnO4- + NO2- + H+  Mn2+ + NO3- + H2O
4.
Balanced the redox equation in base
Cr(OH)3(s) + ClO-  Cl- + CrO42-
5.
Balance the redox equation in acid:
Zn(s) + NO3-  Zn2+ + NH4+.
6.
Balanced the redox equation in base:
NO2- + Al(s)  NH3(aq) + Al(OH)4-
7.
Temperature in K
Balance the disproportionation reaction (same element
undergoes oxidation and reduction): Cl2(g)  Cl- + ClO-.
Moles of H2
in
acid
Mass of Zinc lost
Volume in L
Pressure in atm
Molar Mass of Zinc
%
g.
h.
Why is it necessary to equalize the water level inside
the volumetric flask with the water level in the beaker?
Calculate the average current during the electrolysis
by completing the chart.
in
base
8. Balance the reaction: P4(s)  H2PO2- + PH3(g)
in
acid
in
base
9. Write net ionic equations for the following redox reactions.
a. Solid iron + iron (III) sulfate
Total charge Q in C
Total time T in s
Current I (I = Q/t)
i.
Calculate the approximate voltage needed to perform
the electrolysis. (Assume standard conditions)
b.
Potassium dichromate + acidified hydrogen peroxide.
10. The mass percent of H2O2 in a hydrogen peroxide solution is 13. Using standard reduction potentials, calculate the standard
determined by titration with an acidified solution of KMnO4.
voltage for each of the following reactions.
a. How many moles of MnO4- are needed to react all the
Cl2(g) + 2 I-  2 Cl- + I2(s)
H2O2 if 2.647 mL of 0.0200 M KMnO4 is added?
Ni(s) + 2 Fe3+  Ni2+ + Fe2+
Fe(s) + 2 Fe3+  2 Fe2+
b.
Write the balanced net ionic equation for the reaction.
c.
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
11. Use the Standard Reduction Potential Chart to answer the
following:
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.
Balanced reaction
Eo
Zn  Zn2+
2 Al3+ + 3 Ca(s)  2 Al(s) + 3 Ca2+
14. Using standard reduction potentials, highlight the stronger
reducing agent for each of the following pairs.
Fe(s) or Mg(s) Ca(s) or Al(s) H2(g) or H2S(g) Sn2+ or Fe2+
15. 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 which at
the cathode.
b.
Calculate Eotot for the cell.
c.
Calculate Go.

d.
Write the Nernst equation for the overall reaction.

e.
Calculate the voltage when [H+] = 2.0 x 10-4 M and
[SO42-] = 0.50 M.

f.
What would happen to the voltage of the battery if the
concentration of [H+] increased justify your answer)?

C. Voltaic (Galvanic) Cell
16. Answer the following questions based on the diagram.
Br2  BrOverall
H2  H+
Ag+  Ag
Overall
i. A spontaneous reaction has a (positive/negative) Eotot.
j. Is the redox reaction, Cu(s) + Cl2(g)  CuCl2(s),
spontaneous? Support your answer with calculations.
12. 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.
a. which
species is reduced?
species is oxidized?
species gives up its electron?
species accepts electrons?
ion has a higher electron affinity?
ion passes through the membrane?
species is the reducing agent?
species is the oxidizing agent?
electrode is the cathode?
electrode is the anode?
Zn Zn2+ Cu Cu2+
b.
Which side is the reduction half cell?
c.
Where is the porous membrane?
d.
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?
17. Complete the voltaic cell drawing using the following halfreactions.
Ni2+ + 2 e-  Ni(s)
Eored = -0.25 V
2+
Fe + 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.)
e-

Voltage =
e-

20. 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?
d.
Does the Pb electrode gain or lose mass as the cell
reaction proceeds?
D. Electrolytic Cell
21. Answer the following questions based on the electrolysis of
fused (melted) sodium chloride.
oxidation half reaction
reduction half reaction
overall reaction
18. 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.
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.
Which half-reaction from (b) is higher on the Standard
Reduction Potential Chart?

19. 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.
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?
electrode is the cathode?
electrode is the anode?
b. Write the equations for the electrolysis of NaCl.
oxidation half reaction
Na+
reduction half reaction
overall reaction
c. What mass of Na(l) is produced using a current of
3.00 A for one hour?
Indicate the signs of the two electrodes.

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?
22. Which cell...
a. has a battery?
b. has a salt bridge?
c. is spontaneous?
d. has a positive anode?
Electrolytic
Voltaic
23. Write equations for the oxidation, reduction and overall
reactions for the electrolysis of the salt solutions.
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 K2Cr2O7 solution is added to Na2S solution,
green Cr3+ ions and free S are formed. Which is the best
reducing agent?
(A) K2Cr2O7 (B) Na2S
(C) Cr3+
(D) S
6.
_Ag+ + _AsH3 + _OH-  _Ag + _H3AsO3 + _H2O
When the equation is balanced with lowest whole-number
coefficients, the coefficient for OH- is
(A) 2
(B) 4
(C) 5
(D) 6
7.
_Cr2O72- + _e- + _H+  _Cr3+ + _H2O
When the half reaction is balanced with the lowest wholenumber coefficients, the coefficient for H+ is
(A) 2
(B) 6
(C) 7
(D) 14
8.
_Mg + _NO3– + _H+  _Mg2+ + _NH4+ + _H2O
When the equation is balanced with lowest whole-number
coefficients, the coefficient for H+ is
(A) 1
(B) 3
(C) 5
(D) 10
9.
_CrO2– + _OH–  _CrO42– + _H2O + _e–
When the half-reaction is balanced, what is the ratio of the
coefficients OH–/CrO2–?
(A) 1:1
(B) 2:1
(C) 3:1
(D) 4:1
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
NaCl
CuSO4
Ba(OH)2
HNO3
Na2CO3
KF
24. An electrolytic cell contains a solution of Cr(NO3)3.
a. Write the equations for the anode, cathode and overall
reactions.
anode
cathode
overall
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 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.
(4) Concentration of H+ after 30.0 min. of electrolysis.
(5) Mass of Cr(s) that plate out at the cathode.
11. In which species does sulfur have the same oxidation
number as it does in H2SO4?
(A) H2SO3 (B) S2O32- (C) S2(D) SO2Cl2
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.
12.
2 HClO + 3 O2  2 HClO4
As the reaction represented above 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
When the equation is balanced with lowest whole-number
coefficients, the coefficient for H+ is
(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)
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.
18. 10 HI + 2 KMnO4 + 3 H2SO4  5 I2 + 2 MnSO4 + K2SO4 + 8 H2O
According to the balanced equation above, how many
moles of HI would be necessary to produce 2.5 mol of I2?
(A) 5.0
(B) 8.0
(C) 10.
(D) 12
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)
19. 2 H2O + 4 MnO4- + 3 CIO2-  4 MnO2 + 3 CIO4- + 4 OH(B) (10)(15)(27)/(60)(96,500)
How many moles of ClO2- react with 0.20 L of 0.20 M MnO4-?
(C) (10)(15)(60)(27)/(96,500)(3)
(A) 0.030 (B) 0.053 (C) 0.075 (D) 0.13
(D) (96,500)(27)/(10)(15)(60)(3)
20.
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
Mn2+ + 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
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)
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
Questions 29-33 The spontaneous reaction that occurs when
the cell below operates is 2 Ag+ + Cd  2 Ag + Cd2+
2.
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.
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.
c.
Predict
o.
Justify your prediction.

Calculate the value of Go for the reaction.
30. The silver electrode is made larger.
d.
31. The salt bridge is replaced by a platinum wire.
A current of 1.50 A passes through the cell for 40.0 minutes.
e. Calculate the mass of the Cu(s) that is deposited.
32. Current is allowed to flow for 5 minutes.
f.
33. The silver electrode is replaced by a copper electrode.
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.
3.
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
X3+(aq) + 3 e-  X(s)
Eo = ?
a. Which electrode is the cathode? Justify your answer.
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
b.
What directions do electrons flow in terms of the Sn
and unknown metal strips.
c.
If the cell potential, Eocell, is + 0.60 V, what is the
standard reduction potential for X3+(aq) + 3 e-  X(s)?
Practice Free Response
d.
Identify metal X from the chart of standard potentials.
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?
1.
5 Fe2+ + MnO4- + 8 H+  5 Fe3+ + Mn2+ + 4 H2O
0.500-g of iron(II) compound is dissolved in distilled water,
acidified, and then titrated with 13.5 mL of 0.0200 M KMnO4
to reach the end point. Determine the following
a. The number of moles of MnO4- used.
b.
The number of moles of iron in the sample.
c.
The mass of iron in the sample, in grams.
d.
The mass percent of iron in the compound.