Oxidation and Reduction Reactions Workbook

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Oxidation and Reduction Reactions Workbook
Period/Topic
Worksheets
1.
Oxidation, Reduction, Agents, & Reactions.
2.
Lab: The Strength of Oxidizing Agents.
3.
Oxidation Numbers Spontaneous Reactions
WS 2
4.
Oxidation Numbers, Application to Reactions.
WS 3
5.
Balancing Redox Half Reactions Acid/Base.
WS 4
6.
Balancing Redox Reactions in Acid/Base.
WS 5
7.
Standard Potentials Using Chart.
WS 6
8.
Electrochemical Cells.
WS 7
9.
Electrochemical Cells Lab.
10.
Electrolytic Cells.
11.
Electrolytic Cells Lab.
12.
Application of Electrochemical Cells
13.
Quiz
WS 1
1
2
3
WS 8
4
Application of Electrolytic Cells
WS 9
5
14.
Corrosion, Redox Titrations, Breathalyzer
WS 10
6
15.
Review.
Internet Review
Practice Test 1
16.
Review
Practice Test 2
17.
Test.
Worksheet #1
Redox Half Reactions and Reactions
Define each
1.
2.
3.
4.
Oxidation
Reduction
Oxidizing agent
Reducing agent
Write half reactions for each of the following atoms or ions. Label each as oxidation or reduction.
5.
Al
6.
S
7.
O-2
8.
Ba2+
9.
N3-
10.
Br2
11.
P
12.
Ca
13
Ga3+
14.
S
15.
H2
16.
H+
17.
F-
18.
P3-
Balance each spontaneous redox equation. Identify the entities reduced and oxidized. State the
reducing agent and the oxidizing agent.
19.
Al
&
Zn2+
20.
F2
&
O2-
21.
O2
&
Ca
22.
Al3+
&
Li
Write the oxidation and reduction reactions for each redox reaction. The first one is done for you.
23.
24.
Fe2+
+
⇄
Co
Oxidation:
Co

Co2+
+
2e-
Reduction:
Fe2+
+
2e-

Fe
3 Ag+
Co2+
+
Fe
+
Ni
⇄
Ni3+
+
3 Ag
+
Pb
⇄
Pb2+
+
Cu
+
2 Sn
⇄
O2-
+
2 Sn2+
Oxidation:
Reduction:
25.
Cu2+
Oxidation:
Reduction:
26.
O2
Oxidation:
Reduction:
Co2+
27.
+
2 F-
⇄
Co
+
F2
Oxidation:
Reduction:
28. There are nine formulas for oxidizing agents from questions 19 to 28. List them all. Only
consider formulas that are on the left side of any equation. The first one is done for you.
Zn2+
29. There are nine formulas for reducing agents from questions 19 to 28. List them all. Only
consider formulas that are on the left side of any equation. The first one is done for you.
Al
Worksheet #2
1.
2.
Redox Half Reactions and Reactions
State the Oxidation Number of each of the elements that is underlined.
a) NH3
_____
b) H2SO4
_____
c) ZnSO3
_____
d) Al(OH)3
_____
e) Na
_____
f) Cl2
_____
g) AgNO3
_____
h) ClO4-
_____
i) SO2
_____
j) K2Cr2O4
_____
k) Ca(ClO3)2
_____
l) K2Cr2O7
_____
m) HPO32-
_____
n) HClO
_____
o) MnO2
_____
p) KClO3
_____
q) PbO2
_____
r) PbSO4
_____
s) K2SO4
_____
t) NH4+
_____
u) Na2O2
_____
v) FeO
_____
w) Fe2O3
_____
x) SiO44-
_____
y) NaIO3
_____
z) ClO3-
_____
aa) NO3-
_____
bb) Cr(OH)4
_____
cc) CaH2
_____
dd) Pt(H2O)5(OH)2+ _____
ee) Fe(H2O)63+
_____
ff) CH3COOH
_____
What is the oxidation number of carbon in each of the following substances?
a) CO
_____
b) C
_____
c) CO2
_____
d) CO32-
_____
e) C2H6
_____
f) CH3OH
_____
3.
For each of the following reactants, identify: the oxidizing agent, the reducing agent, the
substance oxidized and the substance reduced.
a)
Cu2+ (aq)
b)
+
→
Zn (s)
Substance oxidized
Oxidizing agent
_____
_____
Cl2 (g) +
→
2 Na (s)
Substance oxidized
Oxidizing agent
Worksheet # 3
Cu(s)
+
Zn2+ (aq)
Substance reduced
Reducing agent
2 Na+ (aq)
_____
_____
+
_____
_____
2 Cl- (aq)
Substance reduced
Reducing agent
____
_____
Spontaneous and Non-spontaneous Redox Reactions
Describe each reaction as spontaneous or non-spontaneous.
→
Fe2+
+
Au
→
Fe2+
+
Pb2+
→
F2
S2O82- + Pb
→
2SO42-
5.
Cu2+
+ 2Br-
→
Cu
6.
Sn2+
+ Br2
→
Sn4+
+ 2Br-
7.
Pb2+
+ Fe2+
→
Fe3+
+ Pb
8.
Can you keep 1 M HCl in an iron container? If the answer is no, write a balanced equation for
the reaction that would occur.
9.
Can you keep 1 M HCl in an Ag container? If the answer is no, write a balanced equation for
the reaction that would occur.
1.
Au3+
2.
Pb
+
3.
Cl2
+
4.
+
Fe3+
Fe3+
F-
+
2Cl+ Pb2+
+ Br2
10.
Can you keep 1 M HNO3 in an Ag container? If the answer is no, write a balanced equation
for the reaction that would occur. (HNO3 consists of two ions H+ and NO3-)
11.
Can you keep 1 M HNO3 in an Au container? If the answer is no, write a balanced equation
for the reaction that would occur. (Remember, HNO3 consists of two ions H+ and NO3-)
12.
Circle each formula that is able to lose an electron
Cl-
O2
13.
Determine the oxidation number for the element underlined.
PbSO4
HP032CaH2
NaIO3
14.
Na+
Fe
Al3+
__________
__________
__________
__________
+
→
Zn
ClO3Na2O2
Al2(SO4)3
C4H12
__________
__________
__________
__________
Al
Zn2+
+
Substance oxidized _______ Oxidizing agent ________
15.
Cr2O72-
+
ClO2- →
Cr3+
+
ClO4-
Substance reduced ________ Oxidizing agent ________
16.
State the Oxidation Number of each of the elements that is underlined.
a) NH3
c) ZnCO3
e) Na
__________
__________
__________
b) H2SO4
d) Al(OH)3
f) Cl2
17. Balance the redox equation using the half reaction method.
Al
&
AgNO3
18. Circle each formula that is able to lose an electron
O2
Cl-
Fe
Na+
__________
__________
__________
Determine the oxidation number for the element underlined.
19.
PbSO4
__________
20.
21.
22.
23.
24.
25.
26.
27.
ClO3HPO32Na202
CaH2
NaIO3
C4H12
Al2(SO4)3
Al3+ +
__________
__________
__________
__________
__________
__________
__________
→
Al
+
__________
Oxidizing agent
Zn
Substance oxidized
28.
Cr2O72- +
ClO2- →
Substance reduced
29.
Cr3+
__________
3As2O3
__________
+ 2H+
Reducing agent
+ 4NO3- + 7H2O + 4 H+ → 6H3AsO4
Substance reduced __________
__________
ClO4-
Oxidizing agent
O3 + H2O + SO2 → SO42- + O2
Substance oxidized__________
30.
+
Zn2+
Reducing agent
__________
+ 4NO
__________
Worksheet # 4
Balancing Redox Reactions
Balance each of the following half-cell reactions. (In each case assume that the reaction takes place
in an ACIDIC solution.) Also, state whether the reaction is oxidation or reduction.
→
1.
S2O32-
SO42-
2.
MnO4-
3.
As
→
AsO43-
4.
Cr3+
→
Cr2O72-
5.
Pb2+
→
PbO2
6.
SO42-
→
S
7.
NO3-
→
NO
8.
NO3-
→
NH4+
9.
BrO3-
→
Br2
→
Mn2+
Balancing Half Cell Reactions
Balance in basic solution.
10.
NO3-
→
NO
11.
MnO4-
→
Mn2+
12.
As
→
AsO43-
13.
Cr3+
→
Cr2O72-
14.
Pb2+
→
PbO2
15.
SO42-
→
S
16.
S2O32-
→
SO42-
17.
NO3-
→
NH4+
18.
BrO3-
→
Br2
19. Determine if each of the following changes is oxidation, reduction or neither.
SO32-
→
SO42-
________________
CaO
→
Ca
________________
CrO42-
→
Cr2O72-
________________
CrO42-
→
Cr3+
________________
20.
2I-
→
I2
________________
IO3-
→
I2
________________
MnO4-
→
Mn2+
________________
ClO2-
→
ClO-
________________
Cr2O72-
+
Fe2+
Substance oxidized
Oxidizing agent
Worksheet # 5
→
Cr3+
_____
_____
+
Fe3+
Substance reduced
Reducing agent
_____
_____
Balancing Redox Reactions in Acid and Basic Solution
Balance each redox equation. Assume all are spontaneous. Use the half reaction method.
1.
O2-
+
F2
2.
Al
+
O2
3.
K
+
Zn+2
Balance each half reaction in basic solution.
4.
Cr2O72 -
→
Cr3+
5.
NO
→
NO3-
6.
SO42-
→
SO2
7.
MnO2
→
Mn2O3
Balance each redox reaction in acid solution using the half reaction method.
8.
H2O2 +
Cr2O72-
→
9.
TeO32- +
N2O4
→
10.
ReO4- +
IO-
→
11.
PbO2 +
I2
→
12.
→
As
O2
Te
+
+
IO3-
NO3-
+
Pb2+
+
IO3-
H2AsO4-
+
AsH3
Balance each redox reaction in basic solution using the half reaction method.
→
13.
O2
+
Cr3+
14.
Te
+
NO3- →
H2O2
+
Cr2O72-
TeO32-
+
N2O4
Cr3+
Re
15.
IO3-
+
Re
→
16.
Pb2+
+
IO3-
→
17.
Cr2O72- +
Hg
→
ReO4-
+
PbO2
Hg2+
IO-
+
+
I2
Cr3+
State of the change represents oxidation, reduction or neither. Use oxidation #s. Remember that if the
oxidation # increases it means oxidation and when it decreases it mean reduction!
18.
MnO2
→
Mn2O3
19.
NH3
→
NO2
20.
HClO4
→
HCl
21.
O2
→
O2-
22.
P2O5
→
P4H10
+
H2O
22.
HS O4-
Determine the oxidation number
23.
H2S O4
24.
P4
23.
NaH
25.
U O3
24.
Na2O 2
26.
U 2O5
25.
P b SO4
Worksheet # 6
1.
Review
Describe each in your own words
a) Oxidation
b) Reduction
c) Oxidizing agent
d) Reducing agent
2. Write half reactions for each. Describe as oxidation or reduction. Circle all oxidizing agents.
a)
b)
c)
d)
e)
f)
Na
Ca
Al3+
F1N2
O2-
3.
Write the reaction between the following: Use the half reaction method.
a)
Ca
+
Al(NO3)3
b)
Sn
+
AgNO3
c)
Sn
+
Au(NO3)3
Cu
Cu+
Al
Al3+
F
O2-
O2
4.
Circle each reducing agent:
5.
Circle each oxidizing agent: F-
6.
Ni+2 reacts with Mn, however, Al+3 does not react with Mn. Rank the oxidizing agents in
order of decreasing strength. Rank the reducing agents in order of decreasing strength.
7.
Ag+ reacts with Pb, however, Ca+2 does not react with Pb. Rank the reducing agents in order
of decreasing strength. Rank the oxidizing agents in order of decreasing strength.
8.
Cl2 reacts with Ag, however, Ag does not react with Mg+2. Rank the oxidizing agents in order
of decreasing strength. Rank the reducing agents in order of decreasing strength.
9.
Ni+2 reacts with Mn, however, Al+3 does not react with Mn. Rank the reducing agents in order
of decreasing strength. Rank the oxidizing agents in order of decreasing strength.
10.
Cl2 reacts with Br-, however, I2 does not react with Br-. Rank the oxidizing agents in order of
decreasing strength. Rank the reducing agents in order of decreasing strength.
Classify as oxidation, reduction or neither.
11.
SO42-
→
S2-
12.
MnO2
→
MnO4-
13.
Cr2O72-
→
CrO42-
14.
IO3-
→
I2
15.
Given the following lab data
SnCl2
Ni(NO3)2
Cr(NO3)3
16.
&
&
&
Ni
Fe
Fe
Spontaneous
Spontaneous
Non spontaneous.
i)
Write three balanced equations.
ii)
Rank the oxidizing agents in decreasing order of strength.
iii)
Rank the reducing agents in decreasing order of strength.
iv)
Will SnCl2 react with Cr? Explain?
v)
Will Fe2+ react with Sn?
Determine the oxidizing and reducing agent. Balance in acidic solution.
MnO4- +
H2S
→
S
+
MnO
17.
Determine the oxidizing and reducing agent. Balance in acidic solution.
SO42- +
Br2
→
S2O32- +
MnO4-
Balance in basic solution
19.
Describe as spontaneous or non-spontaneous. Use your reduction potential chart.
20.
ZnCl2 &
Br2
Cu
&
2+
Fe
H2S
→
18.
a)
c)
+
BrO3-
b)
d)
S
CuCl2 &
H2S
+
MnO
NaCl
&
Al3+
Can you keep HCl in a Zn container? Explain? What about an Au container?
Balance in basic solution
21.
SO42- +
Br2
→
S2O32- +
BrO3-
Classify as an oxidizing agent, reducing agent or both based on its position on the table.
State the Eo or voltage of its position. Some of these are both, so state two voltages and indicate that
it can be an oxidizing and reducing agent.
e.g.
MnO422.
23.
24.
25.
26.
(in acid)
Br2
Fe2+
MnO4- (water)
Ni
Cr3+
oxidizing agent
_________________
_________________
_________________
_________________
_________________
1.51 V
_________________
_________________
_________________
_________________
_________________
27.
H2O
_________________
_________________
Indicate as spontaneous or non-spontaneous.
28.
29.
30.
MnO4- (Alkaline) & Fe2+
HNO3
&
Ag
HCl
&
Mg
Write each oxidation and reduction half reaction for each question above. Determine the Eo for each.
Calculate the Eo for the overall reaction.
34.
35.
36.
Worksheet # 7
Electrochemical Cells
1.
Oxidation is when electrons are
.
2.
Reduction is when electrons are
.
3.
The reducing agent undergoes
.
4.
The oxidizing agent undergoes
.
5.
A negative voltage means the reaction is
6.
In an electrochemical cell electrons exit the electrode which is
7.
In an electrochemical cell the reduction reaction is
oxidation reaction is
.
8.
The cathode is the site of
9.
Anions migrate to the
10.
Anions have a
.
Zn / Zn(NO3)2
and cations migrate to the
charge and cations have a
ll
on the chart, while the
and the anode is the site of
Draw and completely analyze each electrochemical cell.
11.
.
Cu / Cu(NO3)2
charge.
.
.
12.
Ag / AgNO3
ll
H2 / HCl
Worksheet # 8
Electrolytic Cells
1.
In an electrolytic cell, reduction occurs at the
the
electrode.
2.
If there are two possible reduction reactions, the
3.
For reduction, the chart is read from
to
4.
For oxidation, the chart is read from
voltage is
.
to
5.
If there are two possible oxidation reactions, the
6.
Corrosion of a metal is
7.
Electrolysis
8.
Electrochemical cells
9.
Electrolytic cells
10.
What is the standard reference cell?
Molten NaCl
one on the chart occurs.
.
and the sign of the
one on the chart occurs.
.
electrical energy.
electrical energy.
electrical energy.
Draw and completely analyze each electrolytic cell.
11.
electrode and oxidation occurs at
Eo =
v
12.
Aqueous Na2SO4
13.
Liquid K2O
14.
1.0 M LiI
15.
250.0 mL of 0.200 M MnO4- reacts with excess SO3-2. How many grams of
MnO2 are produced? This is Chemistry 11 stoichiometry.
2MnO4- + 3SO32- + H2O → 2MnO2 + 3SO42- + 2OH-
16.
Determine the oxidation number for each underlined atom.
MnO2
Cr2O72IO3C2O42Al(NO3)3
17.
Describe each term:
Salt bridge
Electrolyte
Anode
Cathode
Spontaneous
Electron affinity
18.
What would happen if you used an aluminum spoon to stir a solution of FeSO4(aq)? Write a
reaction and calculate Eo.
19.
Draw an electrochemical cell using Cu and Ag electrodes.
20.
250.0 mL of 0.500 M MnO4- are required to titrate a 100.0 ml sample of SO3-2. Calculate the
[SO3-2]
2MnO4- + 3SO32- + H2O → 2MnO2 + 3SO42- + 2OHHow is the breathalyzer reaction used to determine blood alcohol content (you might need to
look this up in your textbook?
21.
22.
2H+ + Mg → Mg2+ +H2
Oxidizing agent__________
Worksheet # 9
Reducing agent_________
Electrolytic, Electrochemical Cells & Application
Determine the half reactions for each cell and the cell voltage or minimum theoretical voltage and
overall equation.
1.
Ag / Pb electrochemical cell.
Anode:
Anode reaction:
Overall reaction:
2.
Cathode:
Cathode reaction:
Voltage:
ZnCl2(l) electrolytic cell (electrowinning)
Anode:
Anode reaction:
Cathode:
Cathode reaction:
Overall reaction:
3.
CuSO4(aq) electrolytic cell (electrowinning)
Anode:
Anode reaction:
Overall reaction:
4.
Cathode:
Cathode reaction:
MTV:
The electrolysis of 1M NaI (electrowinning)
Anode:
Anode reaction:
Overall reaction:
5.
MTV:
Cathode:
Cathode reaction:
MTV:
The reaction needed to make Al. The electrolyte is
(molten or aqueous).
and its phase is
To lower the mp. from 2000 oC to 800 oC
is used.
Anode:
Cathode:
Anode reaction:
Cathode reaction:
Overall reaction:
6.
The reaction needed to electroplate a copper penny with silver.
Anode:
Anode reaction:
Cathode:
Cathode reaction:
Possible Electrolyte:
7.
The reaction needed to nickel plate a copper penny.
Anode:
Anode reaction:
Cathode:
Cathode reaction:
Possible Electrolyte:
8.
The reaction used in the electrorefining of lead.
Anode:
Anode reaction:
Possible Electrolyte:
Cathode:
Cathode reaction:
Worksheet # 10 Electrolytic, Electrochemical Cells, Corrosion, & Cathodic Protection
Determine the half reactions for each cell and the cell voltage or minimum theoretical voltage.
1.
Zn / Mg electrochemical cell
Anode:
Anode reaction:
Overall reaction:
Cathode:
Cathode reaction:
Voltage:
2.
The electrolytic cell used to produce Al.
Phase (aqueous or molten)
Cathode:
Cathode reaction:
3.
Electrolyte:
Anode:
Anode reaction:
Overall reaction:
The electrolysis KI(aq)
Anode:
Anode reaction:
Overall reaction:
Cathode:
Cathode reaction:
MTV
4.
The electrorefining of Pb
Anode:
Anode reaction:
5.
Cathode:
Cathode reaction:
Nickel plating a iron nail.
Anode:
Cathode:
Anode reaction:
Cathode reaction:
Electrolyte
The -ve side of the power supply is connected to the
6.
Draw an Ag/ Zn electrochemical cell.
7.
Draw a KF(l) electrolytic cell.
8.
Draw a KF(aq) electrolytic cell.
9.
Draw a FeI2(aq) electrolytic cell.
10.
Draw a Cd/Pb electrochemical cell. Cd is not on the reduction chart, however, the Cd
electrode gains mass and the total cell potential is 0.5 v. Determine the half-cell potential for
Cd.
11.
Write the overall reaction and describe the anode and cathode for a Zn/C, fuel, alkaline and
lead/acid cell.
12.
2HIO3 +
5H2SO3
→
oxidizing agent
substance reduced
I2 +
5H2SO4
+
H2O
substance oxidized
reducing agent
13.
What is the electrolyte in a fuel cell?
14.
What is the fuel in a fuel cell?
15.
Describe the differences and similarities between an electrolytic and electrochemical cell.
16.
Describe and give two examples of electrowinning.
17.
Describe and give one example of electrorefining.
18.
List three metals that can be won from aqueous solution.
19.
List three metals that cannot be won from aqueous solution.
20.
List the electrolyte in each of the following.
Fuel cell,
Alkaline battery
Dry Cell (Leclanche)
Lead acid battery
21.
State two metals that can be used to cathodically protect Fe. Describe how they protect iron
from corrosion.
22.
Write the half reaction that describes the corrosion of iron.
23.
Write the half reaction that describes the reduction reaction that occurs when iron corrodes in
air and water.
24.
Why does iron corrode faster in salt water?
25.
Write the anode and cathode reaction in an electrolytic cell with a CaCl2(l) electrolyte.
26.
Explain why you would choose Zn or Cu to cathodically protect iron?
27.
Choose a suitable redox reactant to oxidize Cl- to ClO4- in a redox titration.
28.
Describe as an electrochemical or electrolytic cell:
a) Fuel cell
b) Charging a car battery
c) Discharging a car battery
d) Ni plating
e) Industrial Al production
f) Cl2 production
29.
Write the anode and cathode reactions for each of the above processes.
30.
Al and AgNO3(aq) are mixed and the surface of the Al darkens. List the two oxidizing agents in
decreasing strength. List the two reducing agents in decreasing strength.
31.
Analyze This
Label each anode and cathode.
Write each anode and cathode reaction.
Indicate the ion migration in each cell.
Determine the initial cell voltage of the electrochemical cell.
Determine the MTV for the electrolytic cell.
Will electrolysis occur?
Indicate electron flow.
Indicate all electrodes that gain mass.
Indicate all electrodes that lose mass.
What happens to [NO3-] in the Mg half-cell?
What happens to the [Ag+] in the Ag half-cell?
What happens to [Mg2+] in the Mg half-cell?
What is the equilibrium electrochemical cell potential?
What chemical is made at the Pt electrode on the right?
What chemicals are made at the Pt electrode on the left?
1.0 M KNO3
Mg
Ag
1 M Mg(NO3)2
Pt
1 M AgNO3
Pt
1 M CuSO4
Quiz #1 Agents, Spontaneous Reactions, Oxidation #’s, and Strength
1.
In a redox reaction, the species that loses electrons
A.
is oxidized
B.
C.
D.
2.
Which of the following is the strongest oxidizing agent?
A.
B.
C.
D.
3.
Fe3+ to Fe2+
Fe2+ to MnO4MnO4- to Fe2+
MnO4- to Mn2+
As an element is oxidized, its oxidation number
A.
B.
C.
D.
7.
Fe2+
Fe3+
Mn2+
MnO4-
MnO4- + 5Fe2+ + 8H+ → Mn2+ + 5Fe3+ + 4H2O
During the reaction, electrons transfer from
A.
B.
C.
D.
6.
Ag, Pt, Au
Pt, Au, Ag
Au, Ag, Pt
Au, Pt, Ag
MnO4- + 5Fe2+ + 8H+ → Mn2+ + 5Fe3+ + 4H2O The oxidizing agent in the reaction is
A.
B.
C.
D.
5.
Cu2+
Pb2+
Ni2+
Sn2+
Metallic platinum reacts spontaneously with Au3+(aq) but does not react with Ag+(aq). The
metals, in order of increasing strength as reducing agents, are
A.
B.
C.
D.
4.
is called the cathode
gains mass at the electrode
decreases in oxidation number
increases as electrons are lost
decreases as electrons are lost
increases as electrons are gained
decreases as electrons are gained
A solution of 1.0 M Pb(NO3)2 will not react with a container made of
A.
B.
C.
D.
Cu
Fe
Sn
Zn
8.
A spontaneous redox reaction occurs when a piece of iron is placed in 1.0 M CuSO4. The
reducing agent is
A.
B.
C.
D.
9.
A substance is oxidized when it
A.
B.
C.
D.
10.
Hg2+
weaker reducing agent than Cu2+
weaker oxidizing agent than Cu2+
stronger reducing agent than Cu2+
stronger oxidizing agent than Cu2+
The species which gains electrons in a redox reaction
A.
B.
C.
D.
13.
Ti2+ is a weaker reducing agent than Sn2+
Ti2+ is a weaker oxidizing agent than Sn2+
Ti2+ is a stronger reducing agent than Sn2+
Ti2+ is a stronger oxidizing agent than Sn2+
Consider the following redox reaction : Hg2+ + Cu → Hg + Cu2+ . In this reaction,
A.
B.
C.
D.
12.
loses protons
gains protons
loses electrons
gains electrons
A strip of titanium, Ti, is placed in 1.0 M Sn(NO3)2. The shiny surface of the titanium
darkens, indication that a reaction has occurred. From this observation it may be concluded
that
A.
B.
C.
D.
11.
is a
Fe
Cu2+
H2O
SO42-
loses mass
is oxidized
is the oxidizing agent
increases in oxidization number
Samples of Uranium, Vanadium and Yttrium (U, V, Y) were placed in solutions containing
the metallic ions U3+, V2+, and Y3+. The following observations were recorded.
Trial
1
2
3
4
Ion
U3+
V2+
V2+
Y3+
Metal
Y
U
Y
V
Observation
reaction
reaction
reaction
no reaction
The oxidizing agents from the strongest to the weakest are
A.
B.
C.
D.
V2+, U3+, Y3+
U3+, V2+, Y3+
Y3+, U3+, V2+
V2+, Y3+, U3+
Quiz #2 Agents, Spontaneous Reactions, Oxidation #’s, and Strength
1.
Which of the following pairs of ions will react spontaneously in a solution?
A.
B.
C.
D.
2.
When NO2 reacts to form N2O4 the oxidation number of nitrogen
A.
B.
C.
D.
3.
Cu2+ and Fe2+
Pb2+ and Sn2+
Co2+ and Cr2+
Mn2+ and Cr2+
increases by 2
increases by 4
increases by 8
does not change
Consider the following redox equation:
12H+(aq) + 2IO3-(aq) + 10Fe2+(aq) → 10Fe3+(aq) + I2(s) + 6H2O(l)
The reducing agent is
A.
B.
C.
D.
4.
The oxidation number of nitrogen increases in
A.
B.
C.
D.
5.
I2
H+
Fe2+
IO3-
NO3- → NO
N2O4 → NI3
NH3 → NH4+
NO2 → N2O5
Which of the following represents a balanced reduction half-reaction?
A.
B.
C.
D.
VO2 + 2H+ + 2e- → V2+ + H2O
VO2 + H2 → V2+ + H2O + leVO2 + 2H+ + le- → V2+ + H2O
VO2 + 4H+ + 2e- → V2+ + 2H2O
6.
Consider the following half reaction: Sb2O3 + 6H+ + 6e- ⇄ 2Sb + 3H2O
The oxidation number of antimony in Sb2O3
A.
B.
C.
D.
7.
Consider the following unbalanced half-reaction
The balanced half-reaction would have
A.
B.
C.
D.
8.
10.
HClO2 ⇄ HClO
1 electron on the left
1 electron on the right
2 electrons on the left
2 electrons on the right
The oxidation number of platinum in Pt(H2O)42+ is
A.
B.
C.
D.
9.
increases by 3
increases by 6
decreases by 3
decreases by 6
+2
0
+4
+1/2
Consider the following half-reaction:
BrO- → BrThe balanced equation for the half-reaction is
A.
BrO- + 2H+ + 2e- → Br- + H2O
B.
C.
D.
BrO- + 2H+ → Br- + H2O + 2eBrO- + H2O → Br- + 2OH- + 2eBrO- + H2O + 2e- → Br- + 2OH-
Consider the following redox reaction:
2MnO4- + 5CH3CHO + 6H+ → 5CH3COOH + 2Mn2+ + 3H2O
The species that loses the electron is
A.
B.
C.
D.
11.
H2O
MnO4CH3CHO
CH3COOH
Hydrogen has an oxidation number of –1 in
A.
B.
C.
D.
H2
NaH
H2O
KOH
(basic)
12.
Consider the following:
2NO3- + 4H+ + 2e- → N2O4 + 2H2O
This equation represents
A.
B.
C.
D.
13.
Which of the following half-reactions is balanced?
A.
B.
C.
D.
14.
IO3- + 6H+ +5e- → I2 + 3H2O
IO3- + 6H+ + 4e- →1/2 I2 + 3H2O
IO3- + 6H+ → ½ I2 + 3H2O + 5eIO3- + 6H+ + 5e- → ½ I2 + 3H2O
Consider the following redox reaction: Al + MnO4- + 2H2O →Al(OH)4- + MnO2
The chemical species being oxidized is
A.
B.
C.
D.
15.
reduction
oxidation
neutralization
decomposition
Al
MnO4Al(OH)4MnO2
Consider the following redox reaction:
6H+ + 6I- + ClO3- → 3I2 + 3H2O + ClThe reducing agent is
A.
B.
C.
D.
16.
Nitrogen has an oxidization number of zero in
A.
B.
C.
D.
17.
II2
H+
ClO3-
N2
NO2
NH3
HNO3
When MnO4- reacts to form Mn2+, the manganese in MnO4- is
A.
B.
reduced as its oxidation number increases
reduced as its oxidation number decreases
C.
D.
18.
oxidized as its oxidation number increases
oxidized as its oxidation number decreases
Consider the following reaction:
2HNO3 + 3H2S → 2NO + 3S + 4H2O
The nitrogen in HNO3 undergoes
A.
B.
C.
D.
19.
reduction
oxidation
electrolysis
neutralization
The oxidation number in carbon in CaC2O4 is
A.
B.
C.
D.
+2
+3
+4
+6
20.
Consider the following redox reaction:
2Cr3+(aq) + 3Cl2(aq) + 7H2O(l) → Cr2O72-(aq) + 6Cl-(aq) + 14H+(aq)
The species which loses electrons is
A.
B.
C.
D.
Cl2
Cr3+
H2O
Cr2O72-
Quiz #3 Balancing Redox reactions- Acid & Base Cell Potentials
1.
Consider the following overall reaction:
2Rh+ + Pb(s) → 2Rh(s) + Pb2+The E0 for the half-reaction Rh+ + e- ⇄
A.
B.
C.
D.
2.
Rh is
-0.86 V
-0.60 V
+0.60 V
+0.86 V
Which of the following systems would be correct if the zinc half-cell would have been chosen
as the standard instead of the hydrogen half-cell?
A.
B.
C.
3.
E0 = 0.73 V
The reduction potentials of all the half-cells would remain unchanged
The reduction potentials of all the half-cells would increase by 0.76 V
The reduction potentials of all the half-cells would have positive values
Three beakers contain 1.0 M CuCl2. A piece of metal is placed in each of the beakers
BEAKER
1
2
3
SOLUTION
CuCl2
CuCl2
CuCl2
METAL
Zn
Ag
Ni
Reactions occur in
A.
B.
C.
D.
4.
beaker 2 only
beakers 1, 2, and 3
beakers 1 and 2 only
beakers 1 and 3 only
Consider the following redox reaction:
3SO2 + 3H2O + ClO3- → 3SO42- + 6H+ + Cl-
D.
The reduct
The reduction half-reaction is
A.
B.
C.
D.
5.
What two substances are produced when Cr and 1.0 M MnO4- react in a basic solution?
A.
B.
C.
D.
6.
II2
ClCl2
The substances H2O2, H3PO4 and H2SO3 in order of increasing strengths as oxidizing agents
are.
A
B.
C.
D.
8.
Mn2+ and Cr3+
MnO2 and Cr3+
Mn2+ and Cr2+
MnO2 and CrO42-
Bromine, Br2, will react spontaneously with
A.
B.
C.
D.
7.
ClO3- + 6H+ → Cl- + 3H2O + 6eClO3- + 6H+ + 6e- → Cl - + 3H2O
SO2 + 2H2O → SO42- + 4H+ + 2eSO2 + 2H2O + 2e- → SO42- + 4H+
H2O2, H3PO4, H2SO3
H2SO3, H3PO4, H2O2
H3PO4, H2SO3 , H2O2
H2O2,H2SO3 , H3PO4
Consider the following overall equation for an electrochemical cell:
3Ag+ + Cr → Cr3+ + 3Ag
At standard conditions ,the initial cell voltage is
A.
B.
C.
D.
+0.06 V
+0.39 V
+1.21 V
+1.54 V
9.
A solution of 1.0 M Co(NO3)2 should be stored in a container made of
A.
B.
C.
D.
10.
A strong oxidizing agent has a
A.
B.
C.
D.
11.
Zn
ClSn2+
Fe3+
Which equation represents a redox reaction?
A.
B.
C.
D.
15.
spontaneous and Eo is positive.
spontaneous and Eo is negative.
non-spontaneous and Eo is positive.
non-spontaneous and Eo is negative
Referring to the data booklet, which of the following can act as an oxidizing agent but not as a
reducing agent?
A.
B.
C.
D.
14.
IO3- and I2
SO42- and S
BrO3- and Br AuCl4- and Au
Consider the following redox reaction:
Co2+(aq) + 2Ag(s) → 2Ag+(aq) + Co(s)
The reaction is
A.
B.
C.
D.
13.
weak attraction for electrons
strong attraction for electrons
weak ability to become reduced
strong ability to become oxidized
The two species which react spontaneously in acidic solutions are
A.
B.
C.
D.
12.
tin
zinc
aluminum
magnesium
Pb2+ + 2Cl- → PbCl2
CaO + CO2 → CaCO3
Mg + 2HCl → MgCl2 + H2
HCl + NaOH → NaCl + H2O
In a redox reaction, ClO- was converted to Cl- in a basic solution. The balanced half-reaction
for this process is
A.
B.
ClO- + H2O + 2e- → Cl- + 2OHClO- + 2OH- → Cl- + 2e- + H2O
C.
D.
ClO- + H2O → Cl- + 2e- + 2OHClO- + 2OH- + 2e- → Cl- + H2O
Quiz #4 Electrochemical Cells/Electrolytic Cells
voltmeter
1.0 M KNO3
Zn
Pb
1M Zn(NO3)2
1.
In the electrochemical call above, the electrons flow from
A.
B.
C.
D.
2.
zinc to lead and the mass of zinc increases
zinc to lead and the mass of lead increases
lead to zinc and the mass of zinc increases
lead to zinc and the mass of lead increases
The initial cell voltage is
A.
B.
C.
D.
3.
1M Pb(NO3)2
-0.89 V
-0.63 V
+0.63 V
+0.89 V
In an operating lead-zinc electrochemical cell shown above, the cathode
A.
B.
C.
D.
gains mass as anions are reduced
loses mass as anions are reduced
gains mass as cations are reduced
loses mass as cations are reduced
4.
The equation for the half-reaction at the anode is
A.
B.
C.
D.
5.
Zn2+ + 2e- → Zn
Pb2+ + 2e- → Pb
Zn → Zn2+ + 2ePb → Pb2+ + 2e-
The equation for the half-reaction at the cathode is
A.
B.
C.
D.
6.
Zn2+ + 2e- → Zn
Pb2+ + 2e- → Pb
Zn → Zn2+ + 2ePb → Pb2+ + 2e-
The direction of electron flow in an electrochemical cell is from
A.
B.
C.
D.
7.
anode to cathode through the external wire
cathode to anode through the external wire
anode to cathode through the external wire and back through the salt bridge
cathode to anode through the external wire and back through the salt bridge
Which of the following is formed at the anode during the electrolysis of 1.0 M NaI?
A.
B.
C.
D.
8.
I2
O2
H2
Na
As this cell operates
A.
B.
C.
D.
Cl- is oxidized at the anode
Mg2+ is oxidized at the anode
Cl- is oxidized at the cathode
Mg2+ is oxidized at the cathode
Power Source
Pt
9.
+
Pt
In an operating electrochemical cell, the anions migrate
A.
B.
Molten
2 through the wire
towards
theMgCl
anode
towards the cathode through the wire
C.
D.
towards the anode through the salt bridge
towards the cathode through the salt bridge
voltmeter
1.0 M KNO3
Mg
Cu
1 M Mg(NO3)2
10.
As the above electrochemical cell operates
A.
B.
C.
D.
11.
14.
Cu → Cu2+ + 2eCu2+ + 2e- → Cu
Mg → Mg2+ + 2eMg2+ + 2e- → Mg
In the above electrochemical cell, the initial voltage is
A.
B.
C.
D.
13.
nitrate ions migrate into the copper half-cell
copper(II) ions migrate through the salt bridge
magnesium ions migrate through the salt bridge
potassium ions migrate into the magnesium half-cell
In the above electrochemical cell, the reaction at the anode is
A.
B.
C.
D.
12.
1 M CuSO4
2.03 V
2.52 V
2.71 V
2.89 V
Which of the following aqueous solutions produces H2(g) and O2(g) during electrolysis
A.
1.0 M KI
B.
1.0 M CuI2
C.
1.0 M K2SO4
D.
1.0 M CuSO4
In the electrolysis of molten zinc chloride, the half-reaction at the anode is
A.
B.
C.
D.
Cl2 + 2e- → 2Cl2Cl- → Cl2 + 2eZn2+ 2e- → Zn
Zn → Zn2+ + 2e-
voltmeter
1.0 M KNO3
Ni
Ag
1 M Ni(NO3)2
15.
The initial cell voltage at 25oC is
A.
B.
C.
D.
16.
-1.06 V
-0.54 V
+0.54 V
+1.06 V
The balanced equation for the overall reaction is
A.
B.
C.
D.
17.
1 M AgNO3
Ni+(aq) + Ag(s) → Ag+(aq) + Ni(s)
Ni(s) + Ag+(aq) → Ag(s) + Ni+(aq)
Ni2+(aq) + 2Ag(s) → 2Ag+(aq) + Ni(s)
Ni(s) + 2Ag+(aq) → 2Ag(s) + Ni2+(aq)
This redox reaction occurs because
A.
B.
C.
D.
Ag(s) is a stronger oxidizing agent than Ni(s)
Ag(s) is a weaker reducing agent than Ni(s)
Ag+(aq) is a stronger reducing agent than Ni2+(aq)
Ag+(aq) is a weaker oxidizing agent than Ni2+(aq)
voltmeter
1.0 M KNO3
Au
Pb
18.
The direction of the electron flow is
A.
B.
C.
D.
19.
As the cell operates
A.
B.
C.
D.
20.
-1.37 V
0.00 V
1.37 V
1.63 V
Which of the following is a balanced half-reaction in base?
A.
B.
C.
D.
22.
NO3- and K+ will migrate toward the Pb half-cell
NO3- and K+ will migrate toward the Au half-cell
NO3- migrates toward the Pb half-cell and K+ will migrate toward the Au
NO3- migrates toward the Au half-cell and K+ will migrate toward the Pb
The initial voltage is
A.
B.
C.
D.
21.
from Au to Pb through the wire
from Pb to Au from the wire
from Au to Pb through the salt bridge
from Pb to Au through the salt bridge
Cl2 + 3H2O → ClO3- + 6H+ + 5eCl2 + 6OH- → ClO3- + 5e- + 3H2O
Cl2 + 6H2O → 2ClO3- + 12H+ + 10eCl2 + 12OH- → 2ClO3- + 6H2O + 10e-
In which of the following unbalanced equations does chromium undergo oxidation?
A.
B.
C.
D.
Cr3+ → Cr
Cr3+ → Cr2+
Cr3+ → Cr2O72CrO42- → Cr2O72-
Quiz #5 Application of Cells
1.
The corrosion of iron can be prevented by attaching a piece of zinc to the iron because
A.
B.
iron acts as an anode
zinc reduces more readily than iron
the
C.
D.
2.
An iron spoon is electroplated with copper. The equation representing the reduction
reaction is
A.
B.
C.
D.
A.
B.
C.
D.
4.
ANODE
carbon
pure lead
pure lead
impure lead
3.
In an operating zinc-copper electrochemical
cell, the oxidizing agent
A.
B.
C.
D.
loses electrons at the anode
loses electrons to the cations
gains electrons at the cathode
gains electrons from the anions
extraction of aluminum from bauxite
purification of lead from an impure anode
recovery of zinc from a zinc sulphide solution
production of chlorine from a sodium chloride solution
oxidation of anions
reduction of cations
reduction at the anode
oxidation at the cathode
Hydrogen and oxygen react to provide energy in a
A.
B.
C.
D.
7.
CATHODE
impure lead
carbon
impure lead
pure lead
Electroplating always involves the
A.
B.
C.
D.
6.
Cu2+(aq) + 2e- → Cu(s)
Cu(s) → Cu2+(aq) + 2eFe2+(aq) + 2e- → Fe(s)
Fe(s) → Fe2+(aq) + 2e-
An example of electro refining is the
A.
B.
C.
D.
5.
electrons flow from the zinc to the iron
iron ions form more readily than zinc ions
dry cell
fuel cell
alkaline cell
lead-acid storage cell
En electrolytic process is used to purify impure lead. The electrodes are
8.
In the cell below the half-reaction at the cathode is
Cu2+ + 2e- → Cu(s)
2SO42- → S2O82- + 2eH2O → ½ O2(g) + 2H+ + 2e2H2O + 2e- → H2(g) + 2OH-
A.
B.
C.
D.
Power Source
-
+
Iron Key
Pt
1.0 M CuSO4
9.
In the electrolysis of molten PbBr2, the products at the anode and cathode are
CATHODE
(INERT)
A.
B.
C.
D.
10.
ANODE
(INERT)
Br2
O2
Pb
Br2
H2
Pb
Br2
Pb
Under which conditions could an electrochemical cell provide 0.93V?
Cathode
A.
B.
C.
D.
Anode
Cu
Mg
Mg
Cu
Ag
Pb
11.
The reduction reaction in the above electrochemical cell is
A.
B.
C.
D.
12.
Ba2+
Al3+
Sn2+
Na+
produce fuel
electrolyze fuel
produce hydrogen
produce electricity
If a piece of nickel is to be gold-plated using an electrolytic process, which half-reaction
occurs at the cathode?
A.
B.
C.
D.
17.
the nickel coin must be the cathode
the cathode must be made of copper
the electrons must flow to the anode
the solution must contain nickel ions
The principal function of a fuel cell is to
A.
B.
C.
D.
16.
water forming oxygen gas
water forming hydrogen gas
sea water forming chlorine gas
sea water forming bromine liquid
Which of the following ions can be reduced from an aqueous solution
A.
B.
C.
D.
15.
2e- → Pb
Pb2+ + 2ee- → Ag
Ag+ + e-
To plate a nickel coin with copper
A.
B.
C.
D.
14.
Pb2+ +
Pb →
Ag+ +
Ag →
An industrial process involving electrolysis is the reduction of
A.
B.
C.
D.
13.
Pb
Ag
Ni → Ni2+ + 2eNi2+ + 2e- → Ni
Au → Au3+ + 3eAu3+ + 3e- → Au
Consider the following redox reaction
As2O3 + 2NO3- + 2H2O + 2H+ → 2H3AsO4 + N2O3
In this reaction, nitrogen
A.
B.
C.
D.
18.
In an electrochemical cell, the cathode
A.
B.
C.
D.
19.
Quiz #6
Which of the following metals could be used to cathodically protect a sample of lead?
iron
gold
silver
copper
A piece of iron can be prevented from corroding by
A.
B.
C.
D.
3.
2I- → I2 + 2eNa+ + e- → Na
H2O + ½ O2 + 2H+ + 2e2H2O +2e- → H2 + 2OH-
Corrosion & Cathodic Protection Titration
A.
B.
C.
D.
2.
is reduced
loses mass
is the reducing agent
is the site of reduction
When 1.0 M NaI is electrolyzed, bubbles of gas form on one electrode and a reddish-brown
substance forms on the other. The half-reaction at the cathode is
A.
B.
C.
D.
1.
loses electrons and increases in oxidation number
gains electrons and increases in oxidation number
loses electrons and decreases in oxidation number
gains electrons and decreases in oxidation number
making it a cathode
placing it in an acidic solution
attaching a small piece of lead to it
attaching a small piece of gold to it
To determine the [Fe2+] in a solution of FeSO4 by a redox titration, a suitable reagent would
be an acidified solution of
A.
B.
C.
Cr3+
Mn2+
SO42-
D.
4.
As a metal corrodes,
A.
B.
C.
D.
5.
less than 0.34 V
greater than 1.50 V
greater than 0.85 V but less than 1.50 V
greater than 0.34 V but less than 0.85 V
Consider the following redox equation:
Br2 + SO2 + Na2SO4 + 2H2O → 2H2SO4 + 2NaBr
Which of the following is gaining electrons?
A.
B.
C.
D.
9.
Mn
Cu
Pb
Sn
A student attempted to determine the Eo (volts) of the following half-reaction:
Pd2+ + 2e- → Pd Pd2+ reacts with Cu(s) but not with Hg(l).
Based on the above, the Eo (volts) of a Pd half-cell is
A.
B.
C.
D.
8.
Paint the iron
Cover the iron with grease
Attach a piece of lead tot he iron
Attach a piece of magnesium to the iron
Corrosion of iron can be prevented by attaching a piece of
A.
B.
C.
D.
7.
it gains electrons
it becomes reduced
it acts as a reducing agent
its oxidation number decreases
Which method will cathodically protect a piece of iron?
A.
B.
C.
D.
6.
Cr2O72-
Br2
SO2
H2O
Na2SO4
The reaction that occurs when pieces of lead, zinc, copper and silver are placed in a solution
of Ni(NO3)2 is
A.
B.
C.
D.
Pb + Ni2+ → Pb2+ + Ni
Zn + Ni2+ → Zn2+ + Ni
Cu + Ni2+ → Cu2+ + Ni
2Ag + Ni2+ → 2Ag+ + Ni
voltmeter
1.0 M KNO3
Pb
Cu
1M Pb(NO3)2
10.
In the electrochemical cell above, the electrons flow from
A.
B.
C.
D.
11.
copper to lead through the wire
lead to copper through the wire
copper to lead through the salt bridge
lead to copper through the salt bridge
In the electrochemical cell above, the initial Eo value is
A.
B.
C.
D.
12.
1M Cu(NO3)2
0.03 V
0.21 V
0.29 V
0.47 V
A reaction that occurs during the corrosion of iron is
Fe + 3e- → Fe3+
Fe → Fe2+ + 2eFe2+ + 2e- → Fe
Fe3+ + e- → Fe2+
A.
B.
C.
D.
13.
Consider the following reaction
Zn(s) + 2Ag+(aq) → Zn2+(aq) + 2Ag(s)
What volume of 0.500 M AgNO3 is required to react completely with 6.54 g of zinc?
A.
B.
C.
D.
0.0131 L
0.0262 L
0.200 L
0.400 L
Redox Web Review
1)
Which most readily gains electrons?
Cu
2)
Cu2+
Fe2+2
Zn2+
Au3+
Which most readily loses electrons?
Hg(l)
Cu2+
Sn4+
Ba
Al
Calculate the cell potentials or voltages (E0) Indicate spontaneity.
3.
Cl2 + 2Br- → 2Cl- +Br2
4.
2MnO4- + 5Pb +16H+
5.
Will AgNO3 react with Zn? Write a balanced redox reaction and calculate Eo
6.
→
2Mn2+ + 8H2O + 5Pb2+
What would happen if you used an iron spoon to stir a solution of Al2(SO4)3(aq) ? Write a
balanced redox reaction and calculate Eo.
7.
What are the differences between an electrochemical cell and an electrolytic cell?
Electrochemical cell
8.
Electrolytic cell
What are the similarities between an electrochemical cell and an electrolytic cell?
Electrochemical cell or Electrolytic cell
9.
State how you would determine each of the following in an electrochemical or electrolytic
cell.
Electrochemical Cell
The site of reduction
The site of oxidation
The +ve electrode
The -ve electrode
The anions migrate to the
The cations migrate to the
The electrode that gains mass
The electrode that loses mass
The electrons flow from
Electrolytic Cell
10.
Draw an operating electrochemical cell using an Al half-cell and a Mg half-cell. Label the
parts of the electrochemical cell including the anode or cathode, and all reagents and materials
used. Write the reactions and determine the E0.
11.
Write the half reaction that occurs at each electrode during the electrolysis of aqueous
1.0 M NaI.
Anode :
Cathode :
What is the minimum required voltage for this process?
12.
Write the half reaction that occurs at each electrode during the electrolysis of molten NaI.
Anode :
Cathode :
What is the minimum required voltage for this process?
13.
Aluminum is produced industrially from aluminum oxide, Al2O3. Demonstrate your
understanding of this process by
(i) Describing how the process is carried out,
(ii) Writing equations of the reactions involved in the process, and
(iii) Describing how the problem of the high melting point ofAl2O3 is overcome.
14.
Consider the following redox data:
3V + 2Ga3+ → 3V2+ + 2Ga
Eo = +0.64 V
3V2+ + 2Al
Eo = +0.46 V
→ 3V + 2Al3+
Based on these observations, a student concludes that Ga+3 and Al will react spontaneously.
List the oxidizing agents in order of decreasing strength. Write reduction reactions for each.
Determine the strongest reducing agent. Determine if Ga+3 and Al will react spontaneously.
15.
Balance the equation for the following half reaction occurring in acid solution:
V(s)
16.
HV2O73-
Balance the following redox reaction occurring in basic solution:
MnO4-
17.
→
+
C2O42- →
MnO2
+
CO2
250.0 ml 0.200M MnO4- reacts with excess SO32-. How many grams of MnO2 are produced?
2MnO4- + 3SO32- +H2O → 2MnO2 +3SO42- + 2OH-
18.
Determine the oxidation number for each bold atom.
MnO2
IO3-
Cr2O72-
C2O42-
HOOH
NO3-
H3PO4
Na2C2O4
19.
Al(NO3)3 NH4Cl
I2
N2O3
NaH
Pt(H2O)42+
250.0 mL of 0.500M MnO4- are required to titrate a 100.0 ml sample of SO3-2. Calculate the
[SO3-2]
2MnO4- + 3SO32- +H2O → 2MnO2 + 3SO42-+ 2OH-
20.
How is the breathalyzer reaction used to determine BAC? Write the reaction and describe
how it works.
21.
2H+
+
→
Mg
Mg2+ +
H2
Determine the Oxidizing agent__________ and the Reducing agent_________
22.
Choose a suitable redox reactant to oxidize Cl- toClO4- in a redox titration.
23.
Describe as an electrochemical or electrolytic cell:
a) Fuel cell
b) Charging a car battery
c) Discharging a car battery
d) Ni plating
e) Industrial Al production
f) Cl2 production
g) Electrowinning
24)
Which of the reactants is gaining electrons? Which of the reactants is the oxidizing agent?
Br2 + SO2 + Na2SO4 + H2O → 2H2SO4 + 2NaBr
25)
A student studied the following reactions and she recorded:
Pd2+ + Cu → Pd + Cu2+ spontaneous
Pd2+ + Au → no reaction
Pd2+ + Hg → no reaction
Au3+ + Hg → Au + Hg2+ spontaneous
List the oxidizing agents from strongest to weakest. List the reducing agents from strongest to
weakest. Predict if the reaction will occur.
Au3+
26)
+
Cu
→
Match each type of electrolytic cell with the example cell.
Electrowinning A silver anode oxidizes & Ag reduces on a Cu cathode
Electroplating Pure Pb is reduced at the cathode while impure Pb oxidizes at the anode
Electrorefining Pure Al is reduced at the cathode from molten bauxite (Al2O3).
27.
List the anode, cathode, anode reaction , cathode reaction, and electrolyte for each
commercial electrochemical cell.
Cell
anode anode reaction
cathode cathode reaction
electrolyte
Leclanche or
Common Dry
Cell
Alkaline Cell
Lead Storage
or Car Battery
Fuel Cell
28.
Which of the above cells requires continuous input of O2 and H2 and is produced by Ballard
Industries.
29.
List the anode, cathode, anode reaction, cathode reaction, and electrolyte for each commercial
electrolytic cell.
Cell
anode anode reaction
cathode cathode reaction
electrolyte
Electrolysis of
Molten Al2O3
Electrolysis of
Aqueous NaCl
Silver-plating
a Cu plating
Electrorefining
pure Pb from
impure Pb
30.
Describe each term:
salt bridge
electrolyte
anode
cathode
spontaneous
electron affinity
cation
anion
electrochemical
cell
electrolytic cell
oxidation number
electrolysis
oxidation
reduction
oxidizing agent
reducing agent
electrode
corrosion
electrowinning
electrorefining
over potential
effect
fuel cell
31.
Define corrosion of a metal, and illustrate your definition with reference to an example, using
appropriate equations. Give TWO methods by which corrosion can be prevented and describe
how each method works. The two methods must involve different chemical principles.
32.
Which you would choose Zn or Cu to cathodically protect iron?
33.
A2+ does not react with B, while C2+ reacts with B. Rank the oxidizing agents in decreasing
order of strength. Rank the reducing agents in decreasing order of strength. Will A2+ react
with C?
34.
Write half reactions for each using the reduction table and list the half-cell potential.
Half Reaction
oxidation of water
oxidation of water in acid
reduction of water
reduction of water in alkaline
oxidation of H2 in water
oxidation of H2 in acid
oxidation of H2 in base
reduction of Cr2O72- in acid
reduction of HBr
Eo
35.
Completely analyze the following electrochemical cell.
voltmeter
1.0 M KNO3
Zn
Cu
1 M Zn(NO3)2
1 M Cu(NO3)2
The anode reaction is:
The cathode reaction is:
The electrons flow from ___ to ___
The ions that migrate to the Zn electrode are:
The ions that migrate to the Cu electrode are:
The initial voltage of this cell is:
The voltage of this cell once equilibrium is reached is:
Describe the change in [Cu+2] in the Cu half cell
Describe the change in [NO3-1] in the Zn half cell
36.
Completely analyze the following electrochemical cell.
voltmeter
1.0 M KNO3
H2(g)
Cu
1 M HCl
1 M Cu(NO3)2
The anode reaction is:
The cathode reaction is:
The electrons flow from ___ to ___
The ions that migrate to the Pt electrode are:
The ions that migrate to the Cu electrode are:
The intial voltage of this cell is:
The voltage of this cell once equilibrium is reached is:
Describe the change in [Cu+2] in the Cu half cell
Describe the change in [NO3-1] in the H+/H2 half cell
37.
Completely analyze the following electrolytic cell.
Power Source
-
+
C
C
Molten Al2O3
Anode Reaction
Cathode Reaction
Chemicals produced at the anode
Chemicals produced at the cathode
The electrons flow from __to __
The chemical used to lower the mp is:
Which electrode is the anode ?
38.
Completely analyze the following electrolytic cell. Note that the electrodes are not inert and
because of that, the anode might oxidize.
Power Source
-
+
Cu
Cu
1 M NaF
Anode Reaction
Cathode Reaction
Chemicals produced at the anode
Chemicals produced at the cathode
The electrons flow from
The MTV
Which electrode is the anode ?
Electrochemistry Practice Test # 1
1.
The following represents the process used to produce iron from iron III oxide:
Fe2O3 + 3CO → 2Fe + 3CO2 What is the reducing agent in this process?
A.
B.
C.
D.
Fe
CO
CO2
Fe2O3
2.
Consider the following reaction: 2HNO2 + 2I- + 2H+ → 2NO + I2 +2H2O
The oxidation number for each nitrogen atom
A.
B.
C.
D.
3.
Which of the following reactions is spontaneous?
A.
B.
C.
D.
4.
increases by 1
increases by 2
decreases by 1
decreases by 2
2I- + Ag → Ag+ + I2
Co2+ + Cu → Co + Cu2+
Cu2+ + Pb → Pb2+ + Cu
Ni2+ + 2Ag → 2Ag+ + Ni
Consider the following redox reaction for a lead-acid storage cell:
Pb + PbO2 + 4H+ + 2SO42- → 2PbSO4 + 2H2O
The balanced, reduction half reaction is
A.
B.
C.
D.
5.
-1.12 V
-0.40 V
+0.40 V
+1.12 V
Which of the following involves a nonspontaneous redox reaction?
A.
B.
C.
D.
7.
SO42- → 2PbSO4 + 2e2H+ + SO42- → PbSO4 + 2H2O + 2e+ 4H+ + SO42- + 2e- → PbSO4 + 2H2O
+ 2SO42 + 2H2O + 2e- → PbSO4 + 2OH-
Consider the following reaction: Cd2+(aq) + Zn(s) → Cd(s) Zn2+(aq)
The potential for the reaction is +0.36 V. What is the reduction potential for the cadmium ion?
A.
B.
C.
D.
6.
Pb +
Pb +
PbO2
PbO2
fuel cell
electroplating
redox titration
carbon dry cell
Consider the following redox reaction:
2MnO4- + 16H+ + 5Sn2+ → 2Mn2+ + 8H2O + 5Sn4+
In a redox titration, 0.60 mole of KMnO4 reacts completely with a solution of Sn(NO3)2. How
many moles of Sn(NO3)2 were present in the solution?
A.
0.024 moles
B.
C.
D.
8.
Which of the following is not a redox reaction?
A.
B.
C.
D.
9.
Cathode
O2
Na
Cl2
Cl2
H2
Cl2
H2
Na
1.0 M H+
1.0 M Ag+
1.0 M Sr2+
1.0 M Mn2+
Which of the following half-reactions are balanced?
A.
B.
C.
D.
13.
Anode
A solution containing an unknown cation reacts spontaneously with both zinc and copper. The
unknown cation is
A.
B.
C.
D.
12.
0.26 V
0.28 V
0.54 V
0.80 V
What substances are formed at the anode and cathode during electrolysis of molten sodium
chloride?
A.
B.
C.
D.
11.
Cu + Br2 → CuBr2
CO + H2O → CO2 + H2
CH4 + H2O → CO2 + 2H2O
NaOH + HCl → NaCl + H2O
What is the minimum voltage required to form nickel from an aqueous solution of NiI2 using
inert electrodes?
A.
B.
C.
D.
10.
0.060 moles
1.5 moles
0.30 moles
ClO- + H2O + e- → Cl2 + 2OH2ClO- + H2O + 2e- → Cl2 + 3OH2ClO- + 2H2O + 2e- → Cl2 + 4OH2ClO- + 2H2O → Cl2 + 4OH- + 2e-
Which of the following is a spontaneous redox reaction?
A.
B.
C.
D.
14.
Ag+ + I- → AgI
Ag+ + Fe2+ → Ag + Fe3+
3Ag+ + Au → 3Ag + Au3+
2Ag+ + Ni2+ → 2Ag + Ni
Salting the roads during the winter increases the amount of corrosion of cars. The is because
the salt
A.
B.
C.
D.
reacts with the iron
provides an electrolyte
acts as a reducing agent
acts as an oxidizing agent
Consider the following electrochemical cell for the next five questions.
voltmeter
1.0 M KCl
Cuu
Ni
1M Cu(NO3)2
15.
The half-reaction that occurs at the anode is
A.
B.
C.
D.
16.
0.41 V
0.78 V
0.34 V
0.60 V
The following ions migrate to the Cu electrode
A.
B.
C.
D.
19.
Ni → N2+ + 2eNi2+ + 2e- → Ni
Cu → Cu2+ + 2eCu2+ + 2e- → Cu
The cell potential or Eo is
A.
B.
C.
D.
18.
Ni → N2+ + 2eNi2+ + 2e- → Ni
Cu → Cu2+ + 2eCu2+ + 2e- → Cu
The half-reaction that occurs at the cathode is
A.
B.
C.
D.
17.
1M Ni(NO3)2
K+
Cu2+
ClCl-
Cu2+ Ni2+
Ni2+
NO3NO3- 2e-
The electrons flow
A.
B.
through the salt bridge from Cu to Ni
through the salt bridge from Cu to Ni
C.
D.
20.
Which of the following will not react spontaneously with 1.0 M HCl?
A.
B.
C.
D.
21.
Zn → Zn2+ + 2eZn2+ + 2e- → Zn
2Cl- → Cl2 + 2eCl2 + 2e- → 2Cl-
In the electrolysis of molten zinc chloride, the half-reaction at the anode is
A.
B.
C.
D.
25.
a voltmeter.
a salt bridge.
a power supply.
an aqueous solution.
In the electrolysis of molten ZnCl2 using carbon electrodes, the reaction that occurs at the
anode is
A.
B.
C.
D.
24.
nickel
sodium
aluminum
magnesium
In order for an electrolytic cell to operate, it must have
A.
B.
C.
D.
23.
tin
lithium
mercury
magnesium
Which of the following can be produced by electrolysis from a 1.0 M aqueous solution
containing its ion?
A.
B.
C.
D.
22.
through the wire from Cu to Ni
through the wire from Ni to Cu
Cl2 + 2e- → 2Cl2Cl- → Cl2 + 2eZn2+ + 2e- → Zn
Zn → Zn2+ + 2e-
The corrosion of iron can be prevented by attaching a piece of
A.
B.
C.
D.
Mn
Cu
Pb
Sn
26.
The oxidation number of carbon in CaC2O4 is
A.
B.
C.
D.
27.
+2
+3
+4
+6
To plate a nickel coin with copper,
A.
B.
C.
D.
the nickel coin must be the cathode.
the cathode must be made out of copper
the electrons must flow to the anode
the solution must contain nickel ions
Consider the following electrochemical cell for the next five questions.
voltmeter
1.0 M KNO3
Zn
Cuuu
1M Cu(NO3)2
28.
Which of the following statements apply to this electrochemical cell?
I
II
III
29.
Electrons flow through the wire toward the copper electrode.
The copper electrode increases in mass.
Anions move toward the Zn half-cell.
A.
I and II only
B.
I and III only
C.
II and III only
D.
I, II, and III
The balanced equation for the overall reaction is
A.
B.
C.
D.
30.
1M Zn(NO3)2
Zn + Cu2+ → Cu + Zn2+
Cu + Zn2+ → Zn + Cu2+
Zn2+ + Cu → Cu2+ + Zn
Cu + Zn → Zn + Cu
At equilibrium the voltage of the above cell is
A.
B.
C.
-1.10 V
0.00 V
+0.42 V
D.
31.
This redox reaction occurs because
A.
B.
C.
D.
32.
KCl
NaOH
H2SO4
KOH
The electrolyte used in an automobile battery is
A.
B.
C.
D.
37.
reduced as its oxidation number increases
reduced as its oxidation number decreases
oxidized as its oxidation number increases
oxidized as its oxidation number decreases
The electrolyte used in the alkaline battery is
A.
B.
C.
D.
36.
spontaneous and Eo is positive
spontaneous and Eo is negative
non-spontaneous and Eo is positive
non-spontaneous and Eo is negative
When MnO4- reacts to form Mn2+, the manganese in MnO4- is
A.
B.
C.
D.
35.
-1.10 V
+1.10 V
+0.91 V
+0.86 V
Consider the following redox reaction: Co2+(aq) + 2Ag(s) ⇋ 2Ag+(aq) + Co(s)
The reaction is
A.
B.
C.
D.
34.
Zn is a stronger oxidizing agent than Cu
Zn is a stronger reducing agent than Cu,
Cu is a stronger oxidizing agent than Zn
Zn2+ is a weaker reducing agent than Cu2+
The initial cell voltage at 25 oC is
A.
B.
C.
D.
33.
+1.10 V
KCl
NaOH
H2SO4
KOH
The anode used in the commercial production of Aluminum is
A.
B.
C.
D.
38.
C
Pt
Al
Al2O3
The anode and cathode used in the electrorefining of impure lead to pure lead are
A.
B.
C.
D.
39.
Anode
Cathode
Pure Pb
Impure Pb
Pb2+
Pb
Impure Pb
Pure Pb
Pb
Pb2+
The anode in the LeClanche or common dry cell is
A.
B.
C.
D.
C
Zn
Mg
KOH
40.
Which of the following are electrolytic cells
I
II
III
IV
Electro winning
Electroplating
Charging a car battery
Fuel cell
A.
B.
C.
D.
I and II only
I, II, and III only
II and II only
I, II, III, and IV
Subjective
1.
Balance the following in basic solution.
MnO4-
2.
3.
+
C2O42-
→
MnO2
+
CO2
(basic)
Consider the electrolysis of 1.0 M H2SO4 using platinum electrodes.
a)
Write the oxidation half-reaction
b)
Write the reduction half-reaction
c)
Write the overall reaction and determine the minimum theoretical voltage required.
Consider the following diagram for the electro refining of lead.
Power Source
a)
On the diagram, label the anode and cathode.
b)
Write the formula for a suitable electrolyte
c)
Write the equation for the reduction half-reaction.
4.
Describe two chemically different methods that can be used to prevent corrosion of iron and
explain why each method works.
Method 1:
Explanation:
Method 2:
Explanation:
5.
The data below were obtained in a redox titration of a 25.00 mL sample containing Sn2+ ions
using 0.125 M KMnO4 according to the following reaction:
2MnO4- + 16H+ + 5Sn2+ → 2Mn2+ + 8H2O + 5Sn4+
Calculate the [Sn2+]
Volume of KMnO4 used (mL)
Trial 1
Trial 2
Trial 3
Initial burette reading
Final burette reading
6.
2.00
13.80
13.80
24.55
24.55
35.32
A student wanted to electroplate a coin with copper.
a)
Identify a suitable anode
b)
Identify an appropriate electrolyte
c)
To with battery terminal (positive or negative) should the coin be connected?
7.
Consider the electrolysis of molten magnesium chloride with Cu electrodes (Cu
are not inert and can oxidize: Cl-, or Cu will oxidize)
8.
a)
Identify the product at the anode.
b)
Write the equation for the reduction half-reaction.
c)
Write the equation for the overall reaction.
Completely analyze the following electrochemical cell.
voltmeter
1.0 M KNO3
Snnnn
Mn
electrodes
Chemistry 12
Electrochemistry Practice Test 2
voltmeter
1.0 M KCl
Pd
Ni
1M Pd(NO3)2
1M Ni(NO3)2
1.
As the cell operates, the electrons flow from the nickel electrode to the palladium electrode.
The reaction occurring at the anode is
A
B
C
D
2.
+ 2e+ 2e→ Pb
→ Ni
As the cell operates,
A
B
C
D
3.
Pd → Pd2+
Ni → Ni2+
Pd2+ + 2eNi2+ + 2e-
both the K+ and the NO3- migrate into the nickel half-cell
both the K+ and the NO3- migrate into the palladium half-cell
the K+ migrates into the nickel half-cell and the NO3- migrates into the palladium halfcell
the K+ migrates into the palladium half-cell and the NO3- migrates into the nickel halfcell
The initial cell voltage is 1.21 V. The reduction potential of Pd2+ is
A
B
C
D
-1.21 V
-.95 V
+0.95
+1.21 V
4.
What substances are formed at the anode and cathode during electrolysis of molten
sodium chloride, NaCl(l)?
A
B
C
D
5.
Anode
Cathode
O2
Na
Cl2
Cl2
H2
Cl2
H2
Na
Consider the following electrolytic cell:
Power Source
+
-
Inert
Electrode
Inert
Electrode
Molten NaI(l)
In the cell above
A
B
C
D
6.
I- migrates to the anode and gains electrons
I- migrates to the cathode and loses electrons
Na+ migrates to the anode and loses electrons
Na+ migrates to the cathode and gains electrons
Which of the following are necessary for electroplating to occur using an electrolytic
cell?
I
Two electrodes
II
A metal being reduced
III
A direct current power supply
A
B
C
D
I and II only
I and III only
II and III only
I, II, and III
7.
A fuel cell consumes H2 and O2 gas, uses a KOH electrolyte, and produces electricity.
reaction at the anode is
A
B
2H+ + 2e- → H2
1/2O2 + 2H+ + 2e- → H2O
The
C
D
8.
4OH- → O2 + 2H2O + 4eH2 + 2OH- → 2H2O + 2e-
A student investigating redox reactions recorded the following results:
V2+ + Te2- → no reaction
U4+ + Te2- → U3+ + Te
Based on these results, the strengths of the oxidizing agents, arranged from
weakest, are
A
B
C
D
V2+
U4+
U3+
V2+
Te
Te
Te2Te2-
U4+
V2+
V2+
U3+
9.
What is the minimum voltage required to form nickel from an aqueous solution of
using inert electrodes?
A
B
C
D
0.26 V
0.28 V
0.54 V
0.80 V
strongest to
NiI2
10.
voltmeter
1.0 M KNO3
Zn
Ni
1M Zn(NO3)2
1M Ni(NO3)2
Which of the following occurs as the cell operates?
A
B
C
D
11.
Which of the following reactants would produce an E0 of +0.63 V?
A
B
C
D
12.
Ag+ + I2
Pb2+ + Zn
Mg2+ + Ca
Zn2+ + Mn
The concentration of Fe2+(aq) can be determined by a redox titration using
A
B
C
D
13.
the Zn electrode is reduced and increases in mass
the Zn electrode is reduced and decreases in mass
the Zn electrode is oxidized and increases in mass
the Zn electrode is oxidized and decreases in mass
KBr
SnCl2
KMnO4 (basic)
KBrO3 (acidic)
Which of the following will oxidize Fe2+?
A
B
C
D
I2(s)
Ni(s)
Zn(s)
Br2(l)
14.
The oxidation number of carbon in C2O42- is
A
B
C
D
15.
+3
+4
+5
+6
Consider the following reaction: 3As2O3 + 4NO3- + 7H2O → 6H3AsO4 + 4NO
The oxidizing agent is
A
B
C
D
16.
When W2O5 is converted to WO2 in a redox reaction, the W has been
A
B
C
D
17.
H+
H2O
NO3AsO3
reduced since its oxidation number has increased
reduced since its oxidation number has decreased
oxidized since its oxidation number has increased
oxidized since its oxidation number has decreased
Consider the following:
I
II
III
Water
Oxygen gas
Nitrogen
At 25oC, a piece of iron rusts in the presence of
A
B
C
D
18.
Which of the following represents a redox reaction?
A
B
C
D
19.
I only
III only
I and II only
II and III only
H2CO3 → H2O + CO2
CuS + H2 → H2S + Cu
AgNO3 + NaCl → AgCl + NaNO3
2HCl + Na2SO3 → 2NaCl + H2SO3
The following reaction occurs in an electrochemical cell:
3Cu2+ + Cr → 2Cr3+ + 3Cu
The Eo for the cell is
A
B
C
D
20.
During the corrosion of magnesium, the anode reaction is
A
B
C
D
21.
Zn → Zn2+ + 2e2Cl- → Cl2 + 2eCl2 + 2e- → 2ClZn2+ + 2e- → Zn
Which of the following represents a redox reaction?
A
B
C
D
23.
Mg → Mg2+ + 2eMg2+ + 2e- → Mg
4OH- → O2 + 2H2O + 4eO2 + 2H2O + 4e- → 4OH-
A molten binary salt, ZnCl2, undergoes electrolysis. The cathode reaction is
A
B
C
D
22.
0.40 V
0.75 V
1.08 V
2.50 V
CaCO3 → CaO + CO2
SiCl4 + 2Mg → Si + 2MgCl2
2NaOH + H2SO4 → 2H2O + Na2SO4
AgBr + 2S2O32- → Ag(S2O3)23- + Br-
The process of applying an electric current through a cell to produce a chemical
is called
A
B
C
D
corrosion
ionization
hydrolysis
electrolysis
change
24.
A student investigating redox reactions recorded the following results:
V2+ + Te2- →
no reaction
U4+ + Te2- → U3+ + Te
Based on these results, the strengths of the oxidizing agents, arranged from
weakest, are
A
B
C
D
25.
27.
I2
Cu
H2S
Ag2S
A
B
C
2Cl- → Cl2 + 2e2H+ → H2 + 2eBrO3- + 6H+ + 5e- → ½ Br2 + 3H2O
D
BrO3- + 6H+ → ½ Br2 + 3H2O + 5e-
Which of the following is not a redox reaction?
Cu + Br2 → CuBr2
CO + H2O → CO2 + H2
CH4 + O2 → CO2 + 2H2O
NaOH + HCl → NaCl + H2O
During the electrolysis of 1.0 M Na2SO4, the reaction at the cathode is
A
B
C
D
29.
U4+
V2+
V2+
U3+
Consider the redox reaction: 2BrO3- + 10Cl- + 12H+ → Br2 + 5Cl2 + 6H2O
the oxidation half-reaction ivolved in this reaction is
A
B
C
D
28.
Te
Te
Te2Te2-
A spontaneous redox reaction occurs when Sn2+ is mixed with
A
B
C
D
26.
V2+
U4+
U3+
V2+
Na+ + 1e- → Na
2SO42- → S2O82- + 2e2H2O → O2 + H+ + 4e2H2O + 2e- → H2 + 2OH-
An oxidizing agent will cause which of the following changes?
A
PtO2 → PtO
strongest to
B
C
D
30.
PtO3 → PtO2
Pt(OH)2 → Pt
Pt(OH)22+ → PtO3
Consider the overall reaction of the nickel-cadmium battery:
NiO2(s) + Cd(s) + 2H2O(l) → Ni(OH)2(s) + Cd(OH)2(s)
Which of the following occurs at the anode as the reaction proceeds?
A
B
C
D
Cd loses 2e- and forms Cd(OH)2(s)
Cd gains 2e- and forms Cd(OH)2(s)
NiO2 loses 2e- and forms Ni(OH)2(s)
NiO2 gains 2e- and forms Ni(OH)2(s)
31.
Which of the following can be produced by the electrolysis from a 1.0 M aqueous solution
containing its ions?
A
B
C
D
32.
In the electrolysis of molten ZnCl2 using carbon electrodes, the reaction that occurs at the
anode is
A
B
C
D
33.
nickel
sodium
aluminum
magnesium
Zn → Zn2+ + 2eZn2+ + 2e- → Zn
2Cl- → Cl2 + 2eCl2 + 2e- → 2Cl-
In order for the electrolytic cell to operate, it must have
A
B
C
D
a voltmeter
a salt bridge
a power supply
an aqueous solution
Subjective
1.
a)
Indicate in the blank spaces on the following chart whether or not a reaction will
occur when the metals are added to the aqueous ions.
Pd
Rh
Pt
Pd2+
Rh2+
no reaction
Pt2+
reaction
b)
2.
no reaction
reaction
List the oxidizing agents in order of strongest to weakest
Consider the following reaction for the formation of rust:
Fe(s) + ½ O2(g) + H2O(l) → Fe(OH)2
Describe and explain two methods, using different chemical principles, to prevent
formation of rust.
a)
b)
3.
Consider the following redox reaction:
H2Se + SO42- + 2H+ → Se + H2SO3 + H2O
Calculate the Eo for the reaction.
4.
Balance the following redox reaction in basic solution:
Au
+
Cl-
+
O2
→
AuCl4-
+
OH-
the
5.
Draw and label a simple electrolytic cell capable of electroplating and inert electrode with
silver.
6.
a)
During the production of magnesium metal from seawater, magnesium ions are first
precipitated from seawater as magnesium hydroxide. The magnesium hydroxide is
neutralized by hydrochloric acid, producing magnesium chloride. Write the
neutralization reaction.
b)
The salt produced, magnesium chloride, is dried melted and undergoes electrolysis.
Write the reaction at each electrode.
Anode
Cathode
7.
c)
It is not possible to remove Mg from a 1.0 M solution. Explain why?
d)
Write the anode reaction if Cu electrodes were used instead of C.
Consider the following diagram in the electro refining of lead:
Power Source
Pure Pb
Impure Pb
a)
On the diagram above, label the anode and cathode.
b)
Write the formula for a suitable electrolyte.
c)
Write the equation for the reduction half-reaction.
d)
Write the anode reaction
Oxidation and Reduction Reactions Workbook
Notes- double click on the lesson number and download Power Point Viewer if you do not
have it.
Worksheets
1. Oxidation, Reduction, Agents, & Reactions.
Quiz
WS 1
2. Lab: The Strength of Oxidizing Agents.
3. Oxidation Numbers Spontaneous Reactions
WS 2
4. Oxidation Numbers, Application to Reactions.
1
WS 3
5. Balancing Redox Half Reactions Acid/Base.
WS 4
6. Balancing Redox Reactions in Acid/Base.
WS 5
7. Standard Potentials Using Chart.
WS 6
8. Electrochemical Cells.
WS 7
2
3
9. Electrochemical Cells Lab.
10. Electrolytic Cells.
WS 8
4
12. Application of Electrolytic Cells.
WS 9
5
13. Application of Electrochemical Cells: Bat & Cor.
WS 10
6
14. Breathalyzer and review.
Internet Review
Quizmebc
15. Review
Practice Test # 1
16. Review
Practice Test # 2
11. Electrolytic Cells Lab.
17. Test.
Text book
Hebden
Read Unit V
If you want an A in this class you need to do this!!
Redox Half Reactions and Reactions WS #1
Define each
1. Oxidation
2. Reduction
3. Oxidizing agent
4. Reducing agent
- loss of electrons
- gain of electrons
- causes oxidation by undergoing reduction
- causes reduction by undergoing oxidation
Write half reactions for each of the following atoms or ions. Label each as oxidation or reduction.
Al3+
5.
Al
----------->
6.
S
+
7.
2O2- ---------->
8.
Ba2+
9.
2N3- ---------->
10.
Br2
11.
P +
12.
Ca
----------->
13
Ga3+
+
14.
S
+
15.
H2
--------->
16.
2H+
+
17.
2F-
---------->
F2
18.
P3-
---------->
P +
2e- --------->
+
O2
3e-
+
S2-
reduction
4e-
+
oxidation
2e- -----------> Ba
N2
3e-
reduction
6e-
oxidation
2Br-
reduction
+
2e- --------->
+
P3-
---------->
Ca2+
reduction
2e-
+
3e- -----------> Ga
2e- --------->
2H+
2e- --------->
2e-
H2
+
oxidation
reduction
S2+
oxidation
reduction
oxidation
reduction
2e3e-
oxidation
oxidation
Balance each spontaneous redox equation. Identify the entities reduced and oxidized. State the
reducing agent and the oxidizing agent.
19. Al
&
Zn2+
2Al
+
oxidized
reducing agent
20. F2
&
&
&
3Zn
2O2→
oxidized
reducing agent
4F-
O2
+
Ca
2Ca +
oxidized
reducing agent
22. Al3+
2Al3+ +
O2-
2F2
+
reduced
oxidizing agent
21. O2
3Zn2+
→
reduced
oxidizing agent
O2
→
reduced
oxidizing agent
2Ca2+ +
2O2-
+
3Li+
Li
Al3+ +
reduced
oxidizing agent
3Li
→
oxidized
reducing agent
Al
Label the species that is reduced, that is oxidized, the reducing agent and the oxidizing agent.
Fe2+
23.
Co
Fe
Co
Co2+
+ 2e- oxidation
+
Ni
→
Ni2+
+ 2e-
oxidation
+
Pb
→
Co2+
Fe2+
+
Fe
+
2e-
+
3 Ag
+
1e-
+
Cu
→
reduction
3 Ag+
24.
Ni
Ag
25.
→
→
+
→
Ni3+
Ag+
reduction
Cu2+
Pb2+
→
Pb
Cu
26.
→
+ 2e-
oxidation
+
2 Sn
→
Sn2+
+ 2e-
oxidation
+
2 F-
→
F2
+ 2e-
oxidation
Cu2+
+
2e-
+
2 Sn2+
O2
+
+
F2
Co2+
+
Sn
→ 2O2-
→
O2-
→
Co
→ Coreduction
28. List the species (formulas from above) that lose electrons:
Ni
Pb
Sn
F-
29. List the species (formulas from above) that gain electrons:
Fe2+
Ag+
Cu2+ O2
Co2+
For each of the following reactions, identify:
-The Oxidizing Agent.
-The Reducing Agent.
-The Substance Oxidized.
-The Substance Reduced.
30.
I-
+
4e-
reduction
Co2+
2F-
Co
→
reduction
O2
27.
Pb2+
Cl2
---------->
Substance oxidized
Oxidizing agent
ICl2
Cl-
+
I2
Reducing agent
Substance reduced
ICl2
2e-
31.
Co
+
Fe3+
Substance oxidized
Oxidizing agent
32.
Cr6+
+
Co2+
----------->
Fe2+
Co
Fe3+
Fe2+
Cr6+
Fe2+
Reducing agent
Substance reduced
Co
Fe3+
Cr3+
Fe3+
----------->
Substance oxidized
Oxidizing agent
+
+
Reducing agent
Substance reduced
Fe2+
Cr6+
Redox Half Reactions and Reactions WS #2
1. State the Oxidation Number of each of the elements that is underlined.
a) NH3
c) ZnSO3
e) Na
g) AgNO3
i) SO2
k) Ca(ClO3)2
m) HPO32-
-3
4
0
5
4
5
3
b) H2SO4
d) Al(OH)3
f) Cl2
h) ClO4j) K2Cr2O4
l) K2Cr2O7
n) HClO
6
3
0
7
3
6
1
o) MnO2
q) PbO2
s) K2SO4
u) Na2O2
w) Fe2O3
y) NaIO3
aa) NO3cc) CaH2
ee) Fe(H2O)63+
4
4
6
-1
3
5
5
-1
+3
p) KClO3
r) PbSO4
t) NH4+
v) FeO
x) SiO44z) ClO3bb) Cr(OH)4
dd) Pt(H20)5(0H)2+
ff) CH3COOH
5
2
-3
2
-2
5
4
+3
0
2. What is the oxidation number of carbon in each of the following substances?
a) CO
c) CO2
e) C2H6
2
b) C
2-
4
-3
0
d) CO3
f) CH3OH
4
-2
3. For each of the following reactions, identify: the oxidizing agent, the reducing agent, the
substance oxidized and the substance reduced.
a) Cu2+ (aq)
+
Zn (s)
Substance oxidized
-------->
Zn
Cu (s) +
Zn2+ (aq)
Substance reduced
Cu2+
Cu2+
Oxidizing agent
Reducing agent
Zn
b) Cl2 (g)
+
2 Na (s) -------->
Substance oxidized
Oxidizing agent
2 Cl- (aq)
2 Na+ (aq)
+
Na
Substance reduced
Reducing agent
Cl2
Cl2
Na
WS # 3
Spontaneous and Non-spontaneous Redox Reactions
Describe each reaction as spontaneous or non-spontaneous.
1. Au+3
Fe+3
----->
Fe+3
------>
Fe+2
F-
------>
F2
+ Pb
------>
2SO4-2
+ Pb+2
spontaneous
5.Cu+2
+ 2Br-
------>
Cu
Br2
nonspontaneous
6. Sn+2
+ Br2
------>
Sn+4
+ 2Br-
spontaneous
7. Pb+2
+ Fe+2
------>
Fe+3
+ Pb
nonspontaneous
+
2. Pb
+
3. Cl2
+
4. S2O8-2
Fe+2
+
Au
Pb+2
+
+
+
2Cl-
nonspontaneous (two oxidizing agents)
spontaneous
nonspontaneous
8. Can you keep 1 M HCl in an iron container. If the answer is no, write a balanced equation for the
reaction that would occur. No
Fe
+
2H+
-------->
Fe2+
+
H2
9. Can you keep 1 M HCl in an Ag container. If the answer is no, write a balanced equation for the
reaction that would occur.
Yes. There is no reaction.
10. Can you keep 1 M HNO3 in an Ag container. If the answer is no, write a balanced equation for
the reaction that would occur. (remember HNO3 consists of two ions H+ and NO3-)
No
3Ag
+
NO3- +
4H+
--------> 3Ag+
+
NO
+
2H2O
11. Can you keep 1 M HNO3 in an Au container. If the answer is no, write a balanced equation for
the reaction that would occur. (Remember, HNO3 consists of two ions H+ and NO3-)
Yes. There is no reaction.
12. Circle each formula that is able to lose an elecron
Cl-
O2
Na+
Fe
13. Determine the oxidation number for the element underlined.
14.
15.
PbSO4
HP032-
6
3
ClO3Na2O2
5
-1
CaH2
NaIO3
-1
5
Al2(SO4)3
C4H12
6
-3
Al3+
+
Zn --------->
Substance oxidized
Zn
Cr2O72- + ClO2-
------------>
Substance reduced
Al
+
Zn2+
Oxidizing agent
Cr3+
Cr2O72-
+
Al3+
ClO4-
Oxidizing agent
Cr2O72-
16. State the Oxidation Number of each of the elements that is underlined.
a) NH3
c) ZnCO3
e) Na
-3
4
0
b) H2SO4
d) Al(OH)3
f) Cl2
6
3
0
17. Balance the redox equation using the half reaction method.
Al
+
3Ag+ ---------->
Al3+
+
18. Circle each formula that is able to lose an electron
O2
Cl-
Na+
Fe
Determine the oxidation number for the element underlined.
19.
20.
PbSO4
ClO3-
2
5
3Ag
21.
22.
HPO32Na202
3
-1
23.
24.
25.
26.
CaH2
NaIO3
C4H12
Al2(SO4)3
-1
5
-3
6
27.
Al3+
+
Zn
---------->
Substance oxidized
Cr2O72- +
28.
Al
Zn2+
+
Zn
ClO2- ---------------->
Cr3+
Cr2O72-
Substance reduced
Al3+
Oxidizing agent
+
ClO4-
Oxidizing agent
Cr2O7229.
O3 + H2O + SO2 -----> SO42- +
Substance oxidized
30. 3As2O3
O2
+ 2H+
SO2
Reducing agent
+ 4NO3- + 7H2O + 4 H+ --------> 6H3AsO4
Substance reduced
NO3-
WS # 4
SO2
+ 4NO
Reducing agent
As2O3
Balancing Redox Reactions
Balance each of the following half-cell reactions. (In each case assume that the reaction takes place
in an ACIDIC solution.) Also, state whether the reaction is oxidation or reduction.
1.
5H2O +
S2O32- --------------> 2SO42-
+
10H+ +
8e-
oxidation
2.
8H+
+
5e-
+
MnO4- --------------> Mn2+ +
4H2O
reduction
3.
4H2O +
As
-------------->
AsO43-
+
8H+
+
5e-
Cr2O72-
+
14H+ +
6e-
oxidation
4.
7H2O +
2Cr3+ ----------->
oxidation
5.
Pb2+
2H2O +
--------------> PbO2 +
4H+
+
2e-
oxidation
8H+
6.
+
SO42- +
6e-
--------------> S
NO3- +
3e-
------------->
+
4H2O
reduction
4H+
7.
+
NO
+
2H2O
NH4+ +
3H2O
reduction
10H+ +
8.
8e-
+
NO3- -------------->
10e-
+
2BrO3- -------------->
reduction
12H+ +
9.
Br2
+
reduction
Balancing Half Cell Reactions
Balance in basic solution.
10.
3e-
11.
4H2O +
12.
8OH-
13.
14OH-
14.
4OH-
15.
4H2O +
16.
10 OH-
17.
7H2O +
+
2H2O
5e-
+
As
+
+
+
NO3- --------------> NO
MnO4-
-------------->
AsO43- +
Pb2+ -------------->
6e-
+
4H2O +
8e-
+
8OH-
5e-
+
7H2O +
6e-
PbO2 +
2H2O +
2e-
SO42- --------------> S
S2O32- --------------> 2SO42- +
+
4OH-
--------------> Mn2+ +
2Cr3+ --------------> Cr2O72-
+
+
+
5H2O +
NO3- --------------> NH4+ +
8OH8e-
10 OH-
6H2O
18.
10e-
6H2O +
+
2BrO3- --------------> Br2
+
12 OH-
19. Determine if each of the following changes is oxidation, reduction or neither.
20.
SO32-
-------->
SO42-
oxidation
CaO
-------->
Ca
reduction
CrO42-
-------->
Cr2O72-
neither
CrO42-
-------->
Cr3+
reduction
2I-
-------->
I2
oxidation
IO3-
-------->
I2
reduction
MnO4-
-------->
Mn2+
reduction
ClO2-
-------->
ClO-
reduction
Cr2O72-
Fe2+
+
-------->
Cr3+
+
Fe3+
Substance oxidized
Fe2+
Substance reduced
Cr2O72Oxidizing agent
Cr2O72-
Reducing agent
Fe2+
WS #5
Balancing Redox Reactions in Acid and Basic Solution
Balance each redox equation. Assume all are spontaneous. Use the half reaction method.
1.
2O2-
+
2F2
----------->
O2
+
4F-
2.
4Al
+
3O2
----------->
6O2-
+
4Al3+
3.
2K
+
Zn+2
----------->
Zn
+
2K+
Balance each half reaction in basic solution.
4.
Cr2O72-
+
5.
NO
4OH-
6.
2H2O
7.
2MnO2
+
+
7H2O +
6e-
--------------> 14OH-
------------------>
2H2O
-------------->
SO2
-------------->
Mn2O3
2e-
+
SO42-
+
H2 O
+
2e-
+
2Cr3+
+
NO3- + 3e+
4OH+
2OH-
Balance each redox reaction in acid solution using the half reaction method.
8.
8H+
+
Cr2O72-
3H2O2 +
------->
3O2
2Cr3+ +
+
7H2O
9.
TeO32 - + 2N2O4
10.
4H+
+
4ReO4-+
7IO-
11.
8H+
+
5PbO2 +
I2
12.
12H2O + 8As
+
H2 O
------->
4NO3- +
2H+
-------> 7IO3- +
4Re
+
2H2O
-------> 5Pb2+ +
2IO3- +
4H2O
3H2AsO4-
5AsH3
+
------->
Te
+
+
3H+
Balance each redox reaction in basic solution using the half reaction method.
+
8OH-
+ 2Cr3+
13.
3O2
14.
H2 O
15.
7IO3- +
4OH- + 4Re
16.
8OH- +
5Pb2+ +
17.
7H2O
+ Te +
-------> H2O
4NO3- -------> TeO32- +
-------> 4ReO4- +
+
2OH7IO-
2IO3- -------> 5PbO2 +
+ Cr2O72- + 3Hg
-------> 3Hg2+ +
3H2O2
+
+
2N2O4
+
2H2O
Cr2O72-
I2 + 4H2O
14OH-
+ 2Cr3+
State of the change represents oxidation, reduction or neither (use oxidation #s).
18.
19.
20.
21.
22.
-------->
MnO2
NH3
HClO4
O2
P2O5
Mn2O3
NO2
HCl +
O2P4H10
-------->
------->
-------->
-------->
reduction
oxidation
reduction
reduction
reduction
H2O
Determine the oxidation number
23.
H2S O4
6
22.
HS O4-
6
24.
25.
P4
U O3
0
6
23.
24.
NaH
Na2O 2
-1
-1
26.
U 2O5
5
25.
P b SO4
2
WS #6
Review
1. Describe each in your own words
1. Oxidation
2. Reduction
3. Oxidizing agent
4. Reducing agent
- loss of electrons
- gain of electrons
- causes oxidation by undergoing reduction
- causes reduction by undergoing oxidation
2. Write half reactions for each. Describe as oxidation or reduction. Circle all oxidizing agents.
a)
b)
c)
d)
e)
f)
Na
Ca
Al3+
2F1N2
2O2-
-----------> Na+ +
e2+
-----------> Ca
+
2e+ 3e- -----------> Al
----------> F2 +
2e+
6e- ----------> 2N3----------> O2 +
4e-
oxidation
oxidation
reduction
oxidation
reduction
oxidation
3. Write the reaction between the following: Use the half reaction method.
a) Ca +
3Ca
b) Sn +
Sn +
Al(NO3)3
+
2Al3+
-------------> 2Al
+
AgNO3
2Ag+
-------------> 2Ag
+
Sn2+
3Ca2+
c) Sn +
3Sn
Au(NO3)3
+
2Au3+
-------------> 2Au
+
3Sn2+
4. Circle each reducing agent:
Cu
Cu+
Al
Al3+
5. Circle each oxidizing agent:
F-
F
O2-
O2
6. Ni+2 reacts with Mn, however, Al+3 does not react with Mn. Rank the oxidizing agents in order of
decreasing strength. Rank the reducing agents in order of decreasing strength.
strongest oxidizing agent
Ni2+ +
Mn2+ +
Al3+ +
2e2e3e-
----------->
----------->
----------->
Ni
Mn
Al
strongest reducing agent
7. Ag+ reacts with Pb, however, Ca+2 does not react with Pb. Rank the reducing agents in order of
decreasing strength. Rank the oxidizing agents in order of decreasing strength.
strongest oxidizing agent
Ag+
Pb2+
Ca2+
+
+
+
1e2e2e-
----------->
----------->
----------->
Ag
Pb
Ca
strongest reducing agent
8. Cl2 reacts with Ag, however, Ag does not react with Mg+2. Rank the oxidizing agents in order of
decreasing strength. Rank the reducing agents in order of decreasing strength.
strongest oxidizing agent
Cl2
+
Ag+ +
Mg2+ +
2e1e2e-
-------->
----------->
----------->
2ClAg
Mg
strongest reducing agent
9. Ni+2 reacts with Mn, however, Al+3 does not react with Mn. Rank the reducing agents in order of
decreasing strength. Rank the oxidizing agents in order of decreasing strength.
strongest oxidizing agent
Ni2+ +
Mn2+ +
Al3+ +
2e2e3e-
----------->
----------->
----------->
Ni
Mn
Al
strongest reducing agent
10. Cl2 reacts with Br-, however, I2 does not react with Br-. Rank the oxidizing agents in order of
decreasing strength. Rank the reducing agents in order of decreasing strength.
strongest oxidizing agent
Cl2
Br2
I2
+
+
+
2e2e2e-
-------->
-------->
-------->
Classify as oxidation, reduction or neither.
11. SO4212. MnO2
13. Cr2O72-
-------->
-------->
-------->
S2MnO4CrO42-
reduction
oxidation
neither
2Cl2Br2I-
strongest reducing agent
14. IO3-
-------->
I2
reduction
15. Given the following lab data
SnCl2
Ni(NO3)2
Cr(NO3)3
&
&
&
Ni
Fe
Fe
Spontaneous
Spontaneous
Non spontaneous.
i) Write three balanced equations.
Ni
+
Sn2+ -------------> Ni2+
Fe
+
Ni2+ -------------> Fe2+
Fe
+
Cr3+ <------------- Fe2+
+
+
+
Sn
Ni
Cr
ii) Rank the oxidizing agents in decreasing order of strength.
strongest oxidizing agent
Sn2+
Ni2+
Fe2+
Cr3+
+
+
+
+
2e2e2e3e-
----------->
----------->
----------->
----------->
Sn
Ni
Fe
Cr
strongest reducing agent
iii) Rank the reducing agents in decreasing order of strength. See above.
iv) Will SnCl2 react with Cr? Explain? Yes, because Sn2+ is a stronger oxidizing agent than Cr3+ .
v) Will Fe2+ react with Sn?
2H+
16.
+
No, because Fe2+ is a weaker oxidizing agent than Sn2+
2MnO4-
+
5H2S -------->
oxidizing agent
reducing agent
5S
+ 6H2O +
2MnO
17.
2H+ + 10SO42+
4Br2
oxidizing agent
18.
Balance in basic solution
2MnO419.
5H2S --------> 5S +
2MnO + 4H2O +
2OH-
Describe as spontaneous or non-spontaneous. Use your reduction potential chart.
a)
b)
c)
d)
20.
+
----------> 5S2O32- + 8BrO3- + H2O
reducing agent
ZnCl2
CuCl2
Br2
H2S
&
&
&
&
Cu
NaCl
Fe2+
Al3+
nonspontaneous
nonspontaneous
spontaneous
nonspontaneous
Can you keep HCl in a Zn container?
What about an Au container?
No, Spontaneous reaction.
Yes, nonspontaneous reaction.
Balance in basic solution
21.
H2 O
+
10SO42- + 4Br2
------> 5S2O32- +
2OH- + 8BrO3-
Classify as an oxidizing agent, reducing agent or both based on its position on the table.
State the Eoor voltage of its position. Some of these are both, so state two voltages and indicate that it
can be an oxidizing and reducing agent.
e.g.
MnO4- (in acid)
oxidizing agent
1.51 v
22.
23.
24.
25.
26.
27.
oxidizing agent
oxidizing agent / reducing agent
oxidizing agent
reducing agent
oxidizing agent
oxidizing agent / reducing agent
1.09 v
-0.45 v / 0.77 v
0.60 v
-0.26 v
-0.74 v
-0.40 v / +0.80 v
Br2
Fe2+
MnO4- (water)
Ni
Cr3+
H2O
Indicate as spontaneous or non-spontaneous.
28.
29.
MnO4Cu2+
&
&
Fe2+
Br-
non-spontaneous
non-spontaneous
30.
31.
32.
33.
HNO3
MnO4- (acid)
Ni(s)
HCl
&
&
&
&
Ag
H2O
Al3+
Mg
spontaneous
spontaneous
non-spontaneous
spontaneous
Write each oxidation and reduction half reaction for each question above. Determine the Eo for each.
Calculate the Eo for the overall reaction.
MnO4-
34.
MnO4-
+
+ 2H2O + 3e- -------->
MnO2
3(Fe2+ -----------> Fe3+ + 1e-)
2H2O + 3Fe2+ -----------> 3Fe3+ + MnO2
+
4OH-
+
4OH-
+0.60 v
-0.77 v
-0.17 v
35.
36.
NO3- +
NO3- +
37.
4H+ +3e- ----------->
NO
+
3(Ag ----------> Ag +
1e-)
+
4H
+ 3Ag ----------> NO
+
2H2O
+
2H2O + 3Ag+
+0.96 v
-0.80 v
+0.16 v
38.
39.
Mg
2H+
Mg
+
+
2e------> H2
---------->
Mg2+ +
2e2H+ ---------->
Mg2+ +
WS # 7
H2
0.00 v
2.37 v
2.37 v
Electrochemical Cells
1. Oxidation is when electrons are lost.
2. Reduction is when electrons are gained.
3. The reducing agent undergoes oxidation.
4. The oxidizing agent undergoes reduction.
5. A negative voltage means the reaction is nonspontaneous.
6. In an electrochemical cell electrons exit the electrode, which is negative.
7. In an electrochemical cell the reduction reaction is higher on the chart, while the
oxidation reaction is lower.
.
8. The cathode is the site of reduction and the anode is the site of oxidation.
9. Anions migrate to the anode and cations migrate to the cathode.
10. Anions have a negative charge and cations have a positive charge.
Draw and completely analyze each electrochemical cell.
11. Zn / Zn(NO3)2
║
Cu / Cu(NO3)2
.
2 e-
voltmeter
2 e-
1.0 M KNO3
NO3- K+
Zn
Cu
Zn → Zn2+ + 2eoxidation
anode
0.76 v
loses mass
Zn2+
Cu2+
NO3-
NO3-
1 M Zn(NO3)2
1 M Cu(NO3)2
Cu2+ + Zn → Zn2+ + Cu
12. Ag / AgNO3
║
H2 / HCl
1.10 v
Cu has greater electron affinity
Cu2+ + 2e- → Cu
reduction
cathode
0.34 v
gains mass
2 e-
voltmeter
2 e-
1.0 M KNO3
NO3- K+
H2
H2 → 2H+ + 2eoxidation
anode
0.00 v
Cu
H+
Ag+
Cl-
NO3-
1 M HCl
1 M Ag(NO3)2
2Ag+ + H2 → 2Ag + 2H+
0.80 v
Ag has a greater electron
affinity
2Ag+ + 2e- → 2Ag
reduction
cathode
0.80 v
gains mass
WS # 8
1. In an electrolytic cell, reduction occurs at the negative electrode and oxidation occurs at the
positive electrode.
2. If there are two possible reduction reactions, the highest one on the chart occurs.
3. For reduction, the chart is read from left to right.
4. For oxidation, the chart is read from right to left and the sign of the voltage is changed.
5. If there are two possible oxidation reactions, the lowest one on the chart occurs.
6. Corrosion of a metal is oxidation.
7. Electrolysis uses electrical energy.
8. Electrochemical cells produce electrical energy.
9. Electrolytic cells use electrical energy.
10. What is the standard reference cell? hydrogen Eo = O v
Draw and completely analyze each electrolytic cell.
11. Molten NaCl
Power Source
-
+
Pt
Pt
Na+
Cl-
Cathode: Na+ + 1e- →
2e-1.36 v
Na(s)
Anode: 2Cl- → Cl2 +
-2.71 v
Overall: 2Na+ + 2Cl- → Cl2 + 2Na(s) -4.07 v
MTV = +4.07 v
12. Aqueous Na2SO4
Power Source
-
+
C
C
Na+
SO42H2O
Cathode: 2H2O + 2e- → H2 + 2OH- -0.41 v
+ 1/2O2 + 2e-0.82 v
Anode: H2O
Overall: H2O → H2 + 1/2O2
MTV = +1.23 v
-1.23 v
→ 2H+
13. Liquid K2O
Power Source
-
+
Pt
Pt
K+
O2-
Cathode: K+ + 1e- →
4e?v
K(s)
Anode: 2O2- → O2 +
-2.93 v
Overall: 4K+ + 2O2- → O2 + 4K(s)
-? v
MTV = +? v
14. 1.0 M LiI
Power Source
-
+
Pt
Pt
Li+
I-
Cathode: Cathode: 2H2O + 2e- →
-0.54 v
H2 + 2OH-
Overall: 2H2O + 2I- → I2 + H2 + 2OH-
-0.41 v
-0.95 v
Anode: 2I- → I2 + 2e-
MTV = +0.95 v
15. 250ml of 0.200M MnO4- reacts with excess SO3-2. How many grams of MnO2 are produced? This
is Chemistry 11 stoichiometry.
2MnO4- + 3SO3-2 + H2O -----> 2MnO2 + 3SO4-2 + 2OH0.250L MnO4- x 0.200 mol x 2 mol MnO2
L
2 mol MnO4-
x 86.9g = 4.34g
mol
16. Determine the oxidation number for each underlined atom.
MnO2
4
Cr2O7-2 6
IO35
C2O4-2 3
Al(NO3)3
5
17. Describe each term:
Salt bridge- a u-tube filled with salt solution that allows ions to flow in an electrochemical cell.
Electrolyte- a solution that conducts electricity
Anode- an electrode that is the site of oxidation
Cathode- an electrode that is the site of reduction
Spontaneous- a reaction that occurs naturally and has a positive voltage
Electron affinity- the ability of a metal to attract electrons
18. What would happen if you used an aluminum spoon to stir a solution of FeSO4(aq) ? Write a
reaction and calculate Eo.
2Al
+
3Fe2+ -------> 2Al3+ +
would be a reaction!
3Fe
E0 = 1.21 v
Spontaneous. There
19. Draw an electrochemical cell using Cu and Ag electrodes.
Cathode (+)
Ag
Ag+ + 1e---------> Ag
2Ag+ +
Cu ------>
2Ag
+
0.80v
Anode (-)
Cu
Cu -------> Cu2 + 2e
Cu2+
E0 = 0.46 v
spontaneous
-0.34v
20. 250ml of .500M MnO4- are required to titrate a 100ml sample of SO3-2. Calculate the [SO3-2]
2MnO4- + 3SO3-2 + H2O -----> 2MnO2 + 3SO4-2 + 2OH.250L MnO4- x 0.500 mol x 3 mol SO3-2
L
2MnO40.100L
=
1.88M
21. How is the breathalyzer reaction used to determine blood alcohol content (you might need to look
this up in your textbook)?
The breathalyzer reaction uses a spontaneous redox reaction between acidic Cr2O72- and
ethanol C2H5OH. If alcohol is present in your breath sample, it will react with a solution of
Cr2O72- reducing the orange color as it reacts to form Cr3+, which is green. The drunker you
are, the greater the reduction in orange color, which is measured with a spectrophotometer.
22.
2H+ + Mg-----> Mg+2 +H2
H+
Oxidizing agent
WS #9
Reducing agent
Mg
Electrolytic, Electrochemical Cells & Application
Determine the half reactions for each cell and the cell voltage or minimum theoretical voltage and
overall equation.
1. Ag / Pb electrochemical cell.
Anode:
Anode reaction:
1e- -------> Ag
Overall reaction:
Pb
Pb
Cathode:
Cathode reaction:
--------> Pb2+ + 2e-
Pb + 2Ag+
-----> Pb2+ +
2Ag
Ag
Ag+
+
Voltage: 0.93v
2. ZnCl2(l) electrolytic cell (electro-winning)
Anode:
Anode reaction:
---> Zn
Overall reaction:
C
2Cl- --------> Cl2 + 2e2Cl-
+ Zn2+
-----> Cl2
Cathode:
Cathode reaction:
+
Zn
Zn2+
C
+ 2e- ----
MTV: +2.12 v
3. CuSO4(aq) electrolytic cell (electro-winning)
Anode:
C
Cathode:
C
Anode reaction:
H2O --------> 2H+
2+
Cu
+ 2e -------> Cu
Overall reaction:
H2O + Cu2+
+ 1/2O2 + 2e-----> 2H+
Cathode reaction:
+ 1/2O2 +
Cu
MTV: +0.48 v
4. The electrolysis of 1M NaI (electro-winning)
Anode:
Anode reaction:
----> H2 + 2OHOverall reaction:
C
2I- --------> I2 + 2e2H2O + 2I-
Cathode:
C
Cathode reaction: 2H2O + 2e- ---
-----> H2 + 2OH- + I2
MTV: +0.95 v
5. The reaction needed to make Al. The electrolyte is Al2O3 and its phase is molten (molten or
aqueous).
To lower the mp. from 2000 oC to 800 oC cryolite is used.
Anode:
Anode reaction:
-------> Al
Overall reaction:
6.
+ 4Al3+
-----> 3O2 + 4Al
Ag
Cathode:
Ag-----> Ag+ + e-
penny
Cathode reaction:
Ag+ + e- ----->
penny
Cathode reaction:
Ni2+ + 2e- -----
The reaction needed to nickel plate a copper penny.
Anode:
Anode reaction:
> Ni
Possible Electrolyte
8.
6O2-
Cathode:
C
Cathode reaction: Al3+ + 3e-
4e-
+
The reaction needed to electroplate a copper penny with silver.
Anode:
Anode reaction:
Ag
7.
C
2O2- -------> O2
Ni
Cathode:
Ni-----> Ni+2 + 2eNi(NO3)2
The reaction used in the electrorefining of lead.
Anode:
Anode reaction:
Impure Lead
Pb-----> Pb+2 + 2e-
Cathode:
Cathode reaction:
Pure Lead
Pb2+ + 2e- -----> Pb
WS # 10 Electrolytic, Electrochemical Cells, Corrosion, & Cathodic Protection
Determine the half reactions for each cell and the cell voltage or minimum theoretical voltage.
1.
Zn / Mg electrochemical cell
Anode:
Mg
Anode reaction:
Mg --------> Mg2+ + 2eZn+2 + 2e- -------> Zn
Overall reaction:
Mg + Zn2+ -----> Mg2+ +
2.
Zn
Voltage: 1.61v
The electrolytic cell used to produce Al.
Electrolyte:
Anode:
Anode reaction:
-------> Al
Overall reaction:
3.
Cathode:
Zn
Cathode reaction:
Al2O3
Phase (aqueous or molten)
C
2O2- -------> O2 + 4e6O2-
+ 4Al3+
Molten
Cathode:
C
Cathode reaction: Al3+ + 3e-
-----> 3O2 + 4Al
The electrolysis KI(aq)
Anode:
C
Cathode:
C
Anode reaction:
2I- --------> I2 + 2eCathode reaction: 2H2O
+ 2e- -------> H2 + 2OHOverall reaction:
2H2O + 2I-----> H2 + 2OH- + I2
MTV: +0.95 v
4.
The electrorefining of Pb
Anode:
Anode reaction:
5.
Impure Lead
Pb-----> Pb+2 + 2e-
Cathode:
Cathode reaction:
Pure Lead
Pb2+ + 2e- -----> Pb
Nickel plating an iron nail.
Anode:
Anode reaction:
> Ni
Possible Electrolyte
Ni
Cathode:
+2
Ni-----> Ni + 2eNi(NO3)2
nail
Cathode reaction:
Ni2+ + 2e- -----
The -ve side of the power supply is connected to the nail
6. Draw an Ag/ Zn electrochemical cell.
Anode:
Anode reaction:
1e- -------> Ag
Overall reaction:
Zn
Zn
--------> Zn2+ + 2e-
Zn + 2Ag+
-----> Zn2+ +
Cathode:
Cathode reaction:
2Ag
Ag
Ag+
+
Voltage: 1.56v
7. Draw a KF(l) electrolytic cell.
Anode:
Anode reaction:
Overall reaction:
C
Cathode:
C
+
2F --------> F2 + 2e
Cathode reaction: K + e -------> K
2F- + 2K+
-----> Cl2
+ K
MTV: +5.80v
8. Draw a KF(aq) electrolytic cell.
Anode:
C
Anode reaction:
H2O --------> 2H+ + 1/2O2 + 2e2H2O + 2e- -------> H2 + 2OHOverall reaction:
H2 O
-----> H2 + 1/2O2
Cathode:
Cathode reaction:
C
MTV: +1.23 v
9. Draw a FeI2(aq) electrolytic cell.
Anode:
Anode reaction:
2e- -------> Fe
Overall reaction:
C
2I- --------> I2 + 2eFe2+
+ 2I-
Cathode:
C
2+
Cathode reaction: Fe +
-----> Fe
+
I2
MTV: +0.99 v
10. Draw a Cd/Pb electrochemical cell. Cd is not on the reduction chart, however, the Cd electrode
gains mass and the total cell potential is .5v. Determine the half-cell potential for Cd.
Anode:
Pb
Anode reaction:
Pb --------> Pb2+ + 2eCd+2 + 2e- -------> Zn
x volts
Overall reaction:
Pb + Cd2+ -----> Pb2+ +
0.13 + x = 0.50
0.13v
Cd
Cathode:
Cathode reaction:
Cd
Voltage: 0.50v
x = 0.37v
11. Write the overall reaction and describe the anode and cathode for a dry (Leclanche), fuel, alkaline
and lead/acid cell.
Cell
anode
anode reaction cathode cathode reaction
Leclanche or
Common Dry
Cell
Zn
Zn-->Zn+2 + 2e- C
Mn+4 +1e- -----> Mn+3
NH4Cl and MnO2
Alkaline Cell
Zn
Zn-->Zn+2 + 2e- C
Mn+4 +1e- -----> Mn+3
KOH and MnO2
Lead Storage or
Pb
Car Battery
Pb ---> Pb+2+ 2e- PbO2
PbO2 + HSO4- + 3H+ + 2e- H2SO4
----> PbSO4 + 2H2O
Fuel Cell
H2 + 2OH- --->
2H2O + 2e-
½O2 + H2O +2e-----> 2OH- KOH
C
C
electrolyte
12.
2HIO3
+
5H2S
O3 ----------> I2 + 5H2SO4 + H2O
oxidizing agent
HIO3
substance reduced
HIO3
substance oxidized
reducing agent
H2SO3
H2SO3
13.What is the electrolyte in a fuel cell? KOH
14. What is the fuel in a fuel cell?
H2 and O2
15. Describe the differences and similarities between an electrolytic and electrochemical cell.
Electrolytic
Electrochemical
Uses electricity
Nonspontaneous
Makes chemicals
Inert carbon electrodes
The negative electrode is reduction
Produces electricity
Spontaneous
Uses chemicals
Usually has a salt bridge
The higher metal is reduction
Oxidation occurs at the anode and reduction occurs at the cathode.
Anions migrate to the anode and cations migrate to the cathode.
Electrons go from anode to cathode through the wire.
16. Describe and give two examples of electrowinning.
The electrolysis of water to make H2
and O2.
The electrolysis of Al2O3 to make Al and O2.
17. Describe and give one example of electrorefinning. The electrorefinning of Pb.
18. List three metals that can be won from aqueous solution.
Cu
Fe
Sn
19. List three metals that cannot be won from aqueous solution.
Al
Na
Pb
Au
Ag
Zn
K
Li
Ca
Mg
20. What is the electrolyte in a fuel cell, alkaline battery, Dry Cell (Leclanche) and lead acid battery?
KOH
KOH & MnO2
NH4Cl & MnO2
PbSO4
21. State two metals that can be used to cathodically protect Fe. Describe how they protect iron from
corrosion.
Zn and Mg. When attached to Fe they form an electrochemical cell. Zn or Mg is a stronger
reducing agent (lower on the chart) and is the anode and Fe is the cathode. Since the cathode is
the site of reduction, Fe cannot oxidize or corrode.
22. Write the half reaction that describes the corrosion of iron. Fe --------> Fe2+
+2e-
23. Write the half reaction that describes the reduction reaction that occurs when iron corrodes in air
and water. 2e- + H2O + 1/2O2 ----------> 2OH-
24. Why does iron corrode faster in salt water?
the rate of reaction in an electrochemical cell.
The salt acts like a salt-bridge and increases
25. Write the anode and cathode reaction in an electrolytic cell with a CaCl2 (l) electrolyte.
Cathode: Ca2+ + 2e- ---------> Ca
Anode:
2Cl- ----------> Cl2 +
2e-
26. Explain why you would choose Zn or Cu to cathodically protect iron? Zn. It is a stronger
reducing agent than Fe and it will allow Fe to be the cathode, which cannot corrode.
27. Choose a suitable redox reactant to oxidize Cl- to ClO4- in a redox titration.
MnO4- in acid gives a spontaneous reaction as well as a color change from purple to clear.
28. Describe as an electrochemical or electrolytic cell:
a) Fuel cell
electrolytic
c) Discharging a car battery
electrolytic
e) Industrial Al production
electrolytic
electrochemical
b)Charging a car battery
electrochemical
d) Ni plating
electrolytic
f) Cl2 production
29) Write the anode and cathode reactions.
Cell
anode
anode
reaction
cathode
cathode reaction
electrolyte
Cl2 production C
2Cl- ------> Cl2
C
+ 2e-
Na+ + e- -----> Na
NaCl(l)
Leclanche or
Common Dry
Cell
Zn
Zn-->Zn+2 + 2e- C/MnO2
Mn+4 +1e- -----> Mn+3
NH4Cl and MnO2
Nickel Plating
Ni
Ni-->Ni+2 + 2e-
Metal to be
Ni2+ +2e- -----> Ni
plated
Lead Storage or
Car Battery
Pb
Pb ---> Pb+2+
2e-
PbO2
PbO2 + SO4-2 + 4OH-1 + 2e-----> PbSO4 + 2H2O
H2SO4
Fuel Cell
C
H2 + 2OH- --->
2H2O + 2e-
C
O2 + 2H2O +4e-----> 4OH-
KOH
Ni(NO3)2
30) Al and AgNO3(aq) are mixed and the surface of the Al darkens. List the two oxidizing agents in
decreasing strength. List the two reducing agents in decreasing strength.
Oxidizing Agents
Ag+
Al3+
Reducing Agents
Al
Ag
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