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Oxidation and Reduction Reactions
Period/Topic Worksheets Quiz
1. Oxidation, Reduction, Agents, & Reactions. 1
2. Lab: The Strength of Oxidizing Agents.
3. Oxidation Numbers Spontaneous Reactions 2 1
4. Oxidation Numbers, Application to Reactions. 3
5. Balancing Redox Half Reactions Acid/Base. 4 2
6. Balancing Redox Reactions in Acid/Base. 5
7. Standard Potentials Using Chart. 6 3
8. Electrochemical Cells. 7
9. Electrochemical Cells Lab.
10. Electrolytic Cells. 8 4
11. Electrolytic Cells Lab.
12. Application of Electrochemical Cells
13. Application of Electrolytic Cells 9 5
14. Corrosion, Redox Titrations, Breathalyzer 10 6
15. Review. Internet Review Practice Test 1
16. Review Practice Test 2
17. Test.
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Worksheet #1 Redox Half Reactions and Reactions Define each 1. Oxidation2. Reduction3. Oxidizing agent4. 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-
3
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. Fe2+ + Co ⇄ Co2+ + Fe
Oxidation: Co ® Co2+ + 2e-
Reduction: Fe2+ + 2e- ® Fe
24. 3 Ag+ + Ni ⇄ Ni3+ + 3 Ag
Oxidation:
Reduction:
25. Cu2+ + Pb ⇄ Pb2+ + Cu
Oxidation:
Reduction:
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26. O2 + 2 Sn ⇄ O2-
+ 2 Sn2+
Oxidation:
Reduction:
27. Co2+ + 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 Redox Half Reactions and Reactions 1. 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 _____
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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 _____
2. 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. Use oxidation numbers. a) Cu2+ (aq)+ Zn (s) → Cu(s) + Zn2+ (aq)
Substance oxidized _____ Substance reduced _____Oxidizing agent _____ Reducing agent _____
b) Cl2 (g) + 2 Na (s) → 2 Na+ (aq) + 2 Cl- (aq)
Substance oxidized _____ Substance reduced ____Oxidizing agent _____ Reducing agent _____
Worksheet # 3 Spontaneous and Non-spontaneous Redox Reactions Describe each reaction as spontaneous or non-spontaneous. 1. Au3+ + Fe3+ → Fe2+ + Au
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2. Pb + Fe3+ → Fe2+ + Pb2+ 3. Cl2 + F- → F2 + 2Cl- 4. S2O8
2- + Pb → 2SO42- + Pb2+
5. Cu2+ + 2Br- → Cu + Br2 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.
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
O2 Cl- Fe Na+
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13. Determine the oxidation number for the element underlined.PbSO4 __________ ClO3
- __________HP03
2- __________ Na2O2 __________CaH2 __________ Al2(SO4)3 __________NaIO3 __________ C4H12 __________
14. Al3+ + Zn → 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 __________ b) H2SO4 __________c) ZnCO3 __________ d) Al(OH)3 __________e) Na __________ 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. ClO3
- __________21. HPO3
2- __________22. Na202 __________23. CaH2 __________24. NaIO3 __________25. C4H12 __________26. Al2(SO4)3 __________27. Al3+ + Zn → Al + Zn2+
Substance oxidized __________ Oxidizing agent __________
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28. Cr2O72- + ClO2
- → Cr3+ + ClO4-
Substance reduced __________ Oxidizing agent __________
29. O3 + H2O + SO2 → SO42- + O2 + 2H+
Substance oxidized__________ Reducing agent __________ 30. 3As2O3 + 4NO3
- + 7H2O + 4 H+ → 6H3AsO4 + 4NO
Substance reduced __________ Reducing agent __________ 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. S2O3
2- → SO42-
2. MnO4
- → Mn2+
3. As → AsO4
3-
4. Cr3+ → Cr2O7
2-
5. Pb2+ → PbO2
6. SO4
2- → S
7. NO3- → NO
9
8. NO3- → NH4
+
9. BrO3
- → Br2
Balancing Half Cell Reactions
Balance in basic solution. 10. NO3
- → NO
11. MnO4
- → Mn2+
12. As → AsO4
3-
13. Cr3+ → Cr2O7
2-
14. Pb2+ → PbO2
15. SO4
2- → S
16. S2O3
2- → SO42-
17. NO3
- → NH4+
18. BrO3- → Br2
19. Determine if each of the following changes is oxidation, reduction or neither.
10
SO32- → SO4
2- ________________
CaO → Ca ________________
CrO42- → Cr2O7
2- ________________
CrO42- → Cr3+ ________________
2I- → I2 ________________
IO3- → I2 ________________
MnO4- → Mn2+ ________________
ClO2- → ClO- ________________
20. Cr2O7
2- + Fe2+ → Cr3+ + Fe3+
Substance oxidized _____ Substance reduced _____Oxidizing agent _____ Reducing agent _____
Worksheet # 5 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. Cr2O7
2 - → Cr3+
5. NO → NO3-
6. SO42- → SO2
7. MnO2 → Mn2O3
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Balance each redox reaction in acid solution using the half reaction method.
8. H2O2 + Cr2O72- → O2 + Cr3+
9. TeO3
2- + N2O4 → Te + NO3-
10. ReO4- + IO- → IO3
- + Re
11. PbO2 + I2 → Pb2+ + IO3
-
12. As → H2AsO4
- + AsH3
Balance each redox reaction in basic solution using the half reaction method.
13. O2 + Cr3+ → H2O2 + Cr2O7
2-
14. Te + NO3- → TeO3
2- + N2O4
15. IO3
- + Re → ReO4- + IO-
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16. Pb2+ + IO3
- → PbO2 + I2
17. Cr2O7
2- + Hg → Hg2+ + 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 + H2O 21. O2 → O2-
22. P2O5 → P4H10
Determine the oxidation number 23. H2S O4 22. HS O4
- 24. P 4 23. NaH 25. U O3 24. Na2O 2 26. U 2O5 25. P b SO4
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Worksheet # 6 Review 1. Describe each in your own words
a) Oxidationb) Reductionc) Oxidizing agentd) Reducing agent
2. Write half reactions for each. Describe as oxidation or reduction. Circle all oxidizing agents.
a) Nab) Ca c) Al3+
d) F1-
e) N2
f) 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
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.
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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. SO4
2- → S2-
12. MnO2 → MnO4
-
13. Cr2O7
2- → CrO42-
14. IO3
- → I2
15. Given the following lab data
SnCl2 & Ni SpontaneousNi(NO3)2 & Fe SpontaneousCr(NO3)3 & Fe Non spontaneous.
i) Write three balanced equations.
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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?
16. 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.
SO4
2- + Br2 → S2O32- + BrO3
-
18. Balance in basic solution MnO4- + H2S → S + MnO
16
19. Describe as spontaneous or non-spontaneous. Use your reduction potential chart.
a) ZnCl2 & Cu b) CuCl2 & NaClc) Br2 & Fe2+ d) H2S & Al3+
20. Can you keep HCl in a Zn container? Explain? What about an Au container? Balance in basic solution 21. SO4
2- + 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. MnO4
- (in acid) oxidizing agent 1.51 V22. Br2 _________________ _________________23. Fe2+ _________________ _________________24. MnO4
- (water)_________________ _________________25. Ni _________________ _________________26. Cr3+ _________________ _________________27. H2O _________________ _________________
Indicate as spontaneous or non-spontaneous.
28. MnO4- (Alkaline) & Fe2+
29. HNO3 & Ag 30. 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.
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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 on the chart, while the
oxidation reaction is . 8. The cathode is the site of and the anode is the site of . 9. Anions migrate to the and cations migrate to the . 10. Anions have a charge and cations have a charge. Draw and completely analyze each electrochemical cell. 11. Zn / Zn(NO3)2 ll Cu / Cu(NO3)2 12. Ag / AgNO3 ll H2 / HCl
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Worksheet # 8 Electrolytic Cells 1. In an electrolytic cell, reduction occurs at the electrode and oxidation
occurs at the electrode. 2. If there are two possible reduction reactions, the one on the chart occurs. 3. For reduction, the chart is read from to . 4. For oxidation, the chart is read from to and the sign of the
voltage is . 5. If there are two possible oxidation reactions, the one on the chart occurs. 6. Corrosion of a metal is . 7. Electrolysis electrical energy. 8. Electrochemical cells electrical energy. 9. Electrolytic cells electrical energy. 10. What is the standard reference cell? Eo = v Draw and completely analyze each electrolytic cell. 11. Molten NaCl 12. Aqueous Na2SO4
13. Liquid K2O
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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 + 3SO4
2- + 2OH-
16. Determine the oxidation number for each underlined atom. MnO2 Cr2O7
2- IO3- C2O4
2- Al(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.
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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 + 3SO4
2- + 2OH-
21. How is the breathalyzer reaction used to determine blood alcohol content (you might need to look this up in your textbook?
22. 2H+ + Mg → Mg2+ +H2
Oxidizing agent__________ Reducing agent_________ Worksheet # 9 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: Cathode: Anode reaction: Cathode reaction: Overall reaction: Voltage:
2. ZnCl2(l) electrolytic cell (electrowinning)
Anode: Cathode: Anode reaction: Cathode reaction: Overall reaction: MTV:
3. CuSO4(aq) electrolytic cell (electrowinning)
Anode: Cathode:
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Anode reaction: Cathode reaction: Overall reaction: MTV:
4. The electrolysis of 1M NaI (electrowinning)
Anode: Cathode: Anode reaction: Cathode reaction: Overall reaction: MTV:
5. The reaction needed to make Al. The electrolyte is and its phase is
(molten or aqueous).
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: Cathode: Anode reaction: Cathode reaction: Possible Electrolyte:
7. The reaction needed to nickel plate a copper penny.
Anode: Cathode: Anode reaction: Cathode reaction: Possible Electrolyte:
8. The reaction used in the electrorefining of lead.
Anode: Cathode: Anode reaction: Cathode reaction: Possible Electrolyte:
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: Cathode: Anode reaction: Cathode reaction:
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Overall reaction: Voltage: 2. The electrolytic cell used to produce Al.
Electrolyte: Phase (aqueous or molten) Anode: Cathode: Anode reaction: Cathode reaction: Overall reaction:
3. The electrolysis KI(aq)
Anode: Cathode: Anode reaction: Cathode reaction: Overall reaction: MTV
4. The electrorefining of Pb
Anode: Cathode: Anode reaction: Cathode reaction:
5. 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.
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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 → I2 + 5H2SO4 + H2O
oxidizing agent substance oxidized substance reduced 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.
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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 batteryDry 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?
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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 batteryc) Discharging a car batteryd) Ni platinge) Industrial Al productionf) 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.
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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?
Redox Quiz #1 Agents, Spontaneous Reactions, Oxidation #.
1. In a redox reaction, the species that loses electronsA. is oxidizedB. is called the cathodeC. gains mass at the electrodeD. decreases in oxidation number
2. Which of the following is the strongest oxidizing agent?A. Cu2+
B. Pb2+
1.0 M KNO3
1 M Mg(NO3)2
AgMg
1 M AgNO3
Pt
Pt
1 M CuSO4
27
C. Ni2+
D. Sn2+
3. Metallic platinum reacts spontaneously with Au3+(aq) but does not react with
Ag+(aq).
The metals, in order of increasing strength as reducing agents, areA. Ag, Pt, AuB. Pt, Au, AgC. Au, Ag, PtD. Au, Pt, Ag
4. The oxidizing agent in the reaction below isMnO4
- + 5Fe2+ + 8H+ → Mn2+ + 5Fe3+ + 4H2OA. Fe2+
B. Fe3+
C. Mn2+
D. MnO4-
5. MnO4- + 5Fe2+ + 8H+ → Mn2+ + 5Fe3+ + 4H2O
During the reaction, electrons transfer fromA. Fe3+ to Fe2+
B. Fe2+ to MnO4-
C. MnO4- to Fe2+
D. MnO4- to Mn2+
6. As an element is oxidized, its oxidation numberA. increases as electrons are lostB. decreases as electrons are lostC. increases as electrons are gainedD. decreases as electrons are gained
7. A solution of 1.0 M Pb(NO3)2 will not react with a container made ofA. CuB. FeC. SnD. Zn
8. A spontaneous redox reaction occurs when a piece of iron is placed in 1.0 M CuSO4. The reducing agent isA. FeB. Cu2+
C. H2OD. SO4
2-
9. A substance is oxidized when itA. loses protonsB. gains protonsC. loses electronsD. gains electrons
10. 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 thatA. Ti2+ is a weaker reducing agent than Sn2+
28
B. Ti2+ is a weaker oxidizing agent than Sn2+
C. Ti2+ is a stronger reducing agent than Sn2+
D. Ti2+ is a stronger oxidizing agent than Sn2+
11. Consider the following redox reaction : Hg2+ + Cu → Hg + Cu2+ . In this reaction, Hg2+ is a A. weaker reducing agent than Cu2+
B. weaker oxidizing agent than Cu2+
C. stronger reducing agent than Cu2+
D. stronger oxidizing agent than Cu2+
12. The species which gains electrons in a redox reactionA. loses massB. is oxidizedC. is the oxidizing agentD. increases in oxidization number
13. 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 Ion Metal Observation1 U3+ Y reaction2 V2+ U reaction3 V2+ Y reaction4 Y3+ V no reaction
The oxidizing agents from the strongest to the weakest areA. V2+, U3+, Y3+
B. U3+, V2+, Y3+
C. Y3+, U3+, V2+
D. V2+, Y3+, U3+
14. Use the data below to answer the question.
Mo3+(aq) + Ir(s) ® no observable reaction
Ir3+(aq) + Nd(s) ® Nd3+
(aq) + Ir(s)
Mo3+(aq) + Nd(s) ® Nd3+
(aq) + Mo(s)
From these data it can be predicted that the oxidizing agents, listed from strongest to weakest, are:A. Ir3+ Mo3+ Nd3+
B. Mo3+ Nd3+ Ir3+
C. Ir3+ Nd3+ Mo3+
D. Mo Ir Nd
15. Which process could cause X2-(aq) to change to X+
(aq)?
A. Z(s) ® Z3-(aq)
B. Z(s) ® Z3+(aq)
C. Z2-(aq) ® Z+
(aq)
D. Z3-(aq) ® Z(s)
16. In the reaction, what is the reducing agent?
29
4 Zn(s) + 10 H+(aq) + NO3
-(aq) ® NH4
+(aq) + 4 Zn2+
(aq) + 3 H2O(l)
A. Zn(s)
B. H+(aq)
C. Zn2+(aq)
D. NO3-(aq)
16. A student observed the reactions between four different metals and the solutions of their ions, and then recorded these 'spontaneous' reactions.I. W(s) + X+
(aq) ® W+(aq) + X(s)
II. X(s) + Y+(aq) ® X+
(aq) + Y(s)
III. Y(s) + Z+(aq) ® Y+
(aq) + Z(s)
IV. X(s) + W+(aq) ® X+
(aq) + W(s)
V. X(s) + Z+(aq) ® X+
(aq) + Z(s)
If equation I is correct, which equation did the student record incorrectly?A. IIB. IIIC. IVD. V
17. The reaction below proceeds spontaneously.A2+
(aq) + X(s) ® A(s) + X2+(aq)
Elements A(s) and X(s) respectively, could be
A. Cr2+ and CoB. Pb2+ and CuC. Co2+ and CuD. Ni2+ and Zn
18. A piece of Sn(s) was placed in the following solutions:I. Cu(NO3)2(aq)
II. Hg(NO3)2(aq)
III. AgNO3(aq)
IV. Fe(NO3)2(aq)
A spontaneous reaction will occur in all solutions exceptA. IB. IIC. IIID. IV
19. Theoretically, a reducing agent can be described as a substance thatA. loses electrons and becomes reduced.B. loses electrons and causes reduction.C. gains electrons and causes oxidation.D. gains electrons and becomes reduced.
20. In the reactionCu(s) + 2Ag+
(aq) ® Cu2+(aq) + 2Ag(s)
A. Cu(s) is reduced by the oxidizing agent.B. Ag+
(aq) is the reducing agent and Cu(s) is reduced.C. Cu(s) is the reducing agent and Ag+
(aq) is reduced. D. Cu(s) is the oxidizing agent and Ag+
(aq) is oxidized.
21. Which of the following reactions is non-spontaneous?A. I2(s) + Fe(s) ® 2I-
(aq) + Fe2+(aq)
30
B. Ni2+(aq) + Sn2+
(aq) ® Ni(s) + Sn4+(aq)
C. 2Li(s) + 2H2O(l) ® 2Li+(aq) + 2 OH-
(aq) + H2(g)
D. 2Cl2(g) + 2 H2O(l) ® 4Cl-(aq) + O2(g) + 4H+
(aq)
22. Which of the following best describes the process of oxidation?A. the process in which oxygen is producedB. the process in which electrons are addedC. the process in which the oxidation number decreasesD. the process in which the oxidation number increases
23. What is the oxidation number of N in the mercury (II) compound Hg(NH3)2Cl2
A. -6B. -4C. -3D. +2
24. Which of the following combinations will react spontaneously under standard conditions?A. Ag + Br2
B. Ni + Co2+
C. Zn + Mg2+
D. Au + HNO3
25. Identify the substance that is oxidized in the following equation: Br2 + SO2(g) + K2SO4 + 2H2O ® 2H2SO4 + 2KBrA. Br2
B. SO2
C. H2OD. K2SO4
26. What is the reducing agent in the following equation? Fe2+ + 2I- ® Fe(s) + I2(s)
A. I2
B. I-
C. Fe2+
D. Fe3+
27. What is the oxidation number of C in the C3H5O2- ion ?
A. -1/3B. -2/3C. -1D. -2
28. Which of the following best describes what happens when lead solid is placed in a 1.0 M solution of Cu(NO3)2?A. The solution turns a darker blue.B. No changes are observed.C. Copper solid forms on the lead and the solution changes colour.D. The mass of lead solid increases and the solution does not change colour.
29. Identify the reducing agent in the following equation: Zn + 2MnO2 + H2O ⇌ Zn(OH)2 + 2MnO(OH)
31
A. ZnB. H2OC. MnO2
D. Zn(OH)2
30. Consider the following equation: Co + SO42- + 4H+ ⇌ Co2+ + H2SO3 + H2O
Which statement is correct?A. The sulphur is oxidized and the cobalt is reduced.B. The cobalt is oxidized and the sulphur is reduced.C. The hydrogen is reduced and the cobalt is oxidized.D. The hydrogen is reduced and the oxygen is oxidized.
Redox Quiz #2
1. Which of the following pairs of ions will react spontaneously in a solution?A. Cu2+ and Fe2+
B. Pb2+ and Sn2+
C. Co2+ and Cr2+
D. Mn2+ and Cr2+
2. When NO2 reacts to form N2O4 the oxidation number of nitrogenA. increases by 2B. increases by 4C. increases by 8D. does not change
3. Consider the following redox equation:12H+
(aq) + 2IO3-(aq) + 10Fe2+
(aq) → 10Fe3+(aq) + I2(s) + 6H2O(l)
The reducing agent isA. I2
B. H+
C. Fe2+
D. IO3-
4. The oxidation number of nitrogen increases in
A. NO3- → NO
B. N2O4 → NI3
C. NH3 → NH4+
D. NO2 → N2O5
5. Which of the following represents a balanced reduction half-reaction?A. VO2 + 2H+ + 2e- → V2+ + H2OB. VO2 + H2 → V2+ + H2O + le-
C. VO2 + 2H+ + le- → V2+ + H2OD. VO2 + 4H+ + 2e- → V2+ + 2H2O
6. Consider the following half reaction: Sb2O3 + 6H+ + 6e- ⇄ 2Sb + 3H2O
The oxidation number of antimony in Sb2O3
A. increases by 3B. increases by 6C. decreases by 3D. decreases by 6
7. Consider the following unbalanced half-reactionHClO2 ⇄ HClO
32
The balanced half-reaction would haveA. 1 electron on the leftB. 1 electron on the rightC. 2 electrons on the leftD. 2 electrons on the right
8. The oxidation number of platinum in Pt(H2O)42+ is
A. +2B. 0C. +4D. +1/2
9. Consider the following half-reaction :BrO- → Br- (basic)
The balanced equation for the half-reaction isA. BrO- + 2H+ + 2e- → Br- + H2OB. BrO- + 2H+ → Br- + H2O + 2e-
C. BrO- + H2O → Br- + 2OH- + 2e-
D. BrO- + H2O + 2e- → Br- + 2OH-
10. Consider the following redox reaction:2MnO4
- + 5CH3CHO + 6H+ → 5CH3COOH + 2Mn2+ + 3H2OThe species that loses the electron isA. H2OB. MnO4
-
C. CH3CHOD. CH3COOH
11. Hydrogen has an oxidation number of –1 inA. H2
B. NaHC. H2OD. KOH
12. Consider the following:2NO3
- + 4H+ + 2e- → N2O4 + 2H2OThis equation representsA. reductionB. oxidationC. neutralizationD. decomposition
13. Which of the following half-reactions is balanced?A. IO3
- + 6H+ +5e- → I2 + 3H2OB. IO3
- + 6H+ + 4e- →½ I2 + 3H2OC. IO3
- + 6H+ → ½ I2 + 3H2O + 5e-
D. IO3- + 6H+ + 5e- → ½ I2 + 3H2O
14. Consider the following redox reaction:Al + MnO4
- + 2H2O →Al(OH)4- + MnO2
The chemical species being oxidized isA. AlB. MnO4
-
33
C. Al(OH)4-
D. MnO2
15. Consider the following redox reaction:6H+ + 6I- + ClO3
- → 3I2 + 3H2O + Cl-
The reducing agent isA. I-
B. I2
C. H+
D. ClO3-
16. Nitrogen has an oxidization number of zero inA. N2
B. NO2
C. NH3
D. HNO3
17. When MnO4- reacts to form Mn2+, the manganese in MnO4
- isA. reduced as its oxidation number increasesB. reduced as its oxidation number decreasesC. oxidized as its oxidation number increasesD. oxidized as its oxidation number decreases
18. Consider the following reaction:2HNO3 + 3H2S → 2NO + 3S + 4H2O
The nitrogen in HNO3 undergoesA. reductionB. oxidationC. electrolysisD. neutralization
19. The oxidation number in carbon in CaC2O4 isA. +2B. +3C. +4D. +6
20. Consider the following redox reaction:2Cr3+
(aq) + 3Cl2(aq) + 7H2O(l) → Cr2O72-
(aq) + 6Cl-(aq) + 14H+
(aq)
The species which loses electrons isA. Cl2
B. Cr3+
C. H2OD. Cr2O7
2-
21. Which equation represents a redox reaction?A. Pb2+ + 2Cl- → PbCl2
B. CaO + CO2 → CaCO3
C. Mg + 2HCl → MgCl2 + H2
D. HCl + NaOH → NaCl + H2O
22. In a redox reaction, ClO- was converted to Cl- in a basic solution. The balanced half-reaction for this process isA. ClO- + H2O + 2e- → Cl- + 2OH-
B. ClO- + 2OH- → Cl- + 2e- + H2OC. ClO- + H2O → Cl- + 2e- + 2OH-
34
D. ClO- + 2OH- + 2e- → Cl- + H2O
23. Which of the following equations is not predicted to represent a redox reaction?A. 2 H2O(l) + O2(g) ® 2 H2O2(l)
B. 2 Sn2+(aq) ® Sn(s) + Sn4+
(aq)
C. Ag+(aq) + Cl-
(aq) ® AgCl(s)
D. C2H4(g) + 3 O2(g) ® 2 CO2(g) + 2 H2O(g)
24. The metal molybdenum, Mo, can react to produce MoO2(s). The half reaction that explains the change in oxidation state of molybdenum can be writtenA. Mo(s) + 2e- → Mo2+ B. Mo(s) → Mo2+ + 2e-
C. Mo4+ + 4e- → Mo(s)
D. Mo(s) → Mo4+ + 4e-
25. Which incomplete half-reaction is an oxidation?A. 2H+ + 12 O2(g) → H2O(l)
B. Cr2O72-
+ 14H+ → 2 Cr3+
+ 7 H2O(l)
C. K+ → K(s)
D. 2 I- → I2(s)
26. Which of these would react spontaneously with Co(s) but not with H2(g)?A. H2O(l)
B. Ni2+(aq)
C. Ca2+(aq)
D. Br-(aq)
27. The chemical that will reduce copper (II) ions and also oxidize metallic nickel is aqueousA. iron (III) nitrateB. tin (II) sulfateC. iron (II) nitrateD. zinc sulfate
28. F2(g) is a strong reducing agent F2(g) has a strong attraction for electrons. Based on this information, one should determine thatA. the statement and the explanation are true, and that the explanation is
correct for the statement.B. both the statement and the explanation are true, but the explanation is not
correct for the statement.C. the statement is true, but the explanation is false.D. the statement is false, but the explanation is true
29. Solutions containing Fe2+(aq), Sn2+
(aq) and Cr2+(aq) are mixed together. The reaction
most likely to occur isA. 2Fe2+
(aq) + 2 H2O(l) ® Fe(s) + O2(g) + 4H+(aq)
B. Sn2+(aq) + 2Cr2+
(aq) ® Sn(s) + 2 Cr3+(aq)
C. Sn2+(aq) + Cr2+
(aq) ® Sn4+(aq) + Cr(s)
D. Sn2+(aq) + 2 Fe2+
(aq) ® Sn(s) + 2 Fe3+(aq)
30. Four reducing agents listed in order of decreasing strength are W, Z, Y and X. Four statements about the reaction between the reducing agents and their respective oxidizing agents are:I. W(s) + X2+
(aq) ® W2+(aq) + X(s)
35
II. Y(s) + X2+(aq) ® Y2+
(aq) + X(s)
III. W(s) + Z2+(aq) ® no reaction
IV. Y(s) + Z2+(aq) ® Y2+
(aq) + Z(s)
The statement(s) inconsistent with the correct order of reducing agents is (are)A. IV onlyB. III onlyC. I and IID. III and IV
Redox Quiz #3 Balancing Redox reactions- Acid & Base Cell Potentials
1. Consider the following overall reaction:2Rh+
+ Pb(s) → 2Rh(s) + Pb2+- E0 = 0.73 V The E0 for the half-reaction Rh+
+ e- ⇄ Rh isA. -0.86 VB. -0.60 VC. +0.60 VD. +0.86 V
2. 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. The reduction potentials of all the half-cells would remain unchangedB. The reduction potentials of all the half-cells would increase by 0.76 VC. The reduction potentials of all the half-cells would have positive valuesD. The reduction potentials of the hydrogen half-cell decrease by 0.76 V
3. Three beakers contain 1.0 M CuCl2. A piece of metal is placed in each of the beakers
BEAKER SOLUTION METAL1 CuCl2 Zn2 CuCl2 Ag3 CuCl2 Ni
Reactions occur inA beaker 2 onlyB. beakers 1, 2, and 3C. beakers 1 and 2 onlyD. beakers 1 and 3 only
4. Consider the following redox reaction:3SO2 + 3H2O + ClO3
- → 3SO42- + 6H+ + Cl-
The reduction half-reaction isA. ClO3
- + 6H+ → Cl- + 3H2O + 6e-
B. ClO3- + 6H+ + 6e- → Cl - + 3H2O
C. SO2 + 2H2O → SO42- + 4H+ + 2e-
D. SO2 + 2H2O + 2e- → SO42- + 4H+
5. What two substances are produced when Cr and 1.0 M MnO4- react in a basic solution?
A. Mn2+ and Cr3+
B. MnO2 and Cr3+
36
C. Mn2+ and Cr2+
D. MnO2 and CrO42-
6. Bromine, Br2, will react spontaneously withA. I-
B. I2
C. Cl-
D. Cl2
7. The substances H2O2, H3PO4 and H2SO3 in order of increasing strengths as oxidizing agents are.A H2O2, H3PO4, H2SO3 B. H2SO3, H3PO4, H2O2 C. H3PO4, H2SO3 , H2O2
D. H2O2,H2SO3 , H3PO4
8. Consider the following overall equation for an electrochemical cell:3Ag+ + Cr → Cr3+ + 3Ag
At standard conditions ,the initial cell voltage isA. +0.06 VB. +0.39 VC. +1.21 VD. +1.54
9. A solution of 1.0 M Co(NO3)2 should be stored in a container made ofA. tinB. zincC. aluminumD. magnesium
10. A strong oxidizing agent has a A. weak attraction for electronsB. strong attraction for electronsC. weak ability to become reducedD. strong ability to become oxidized
11. The two species which react spontaneously in acidic solutions areA. IO3
- and I2
B. SO42- and S
C. BrO3- and Br -
D. AuCl4- and Au
12. Consider the following redox reaction:Co2+
(aq) + 2Ag(s) → 2Ag+(aq) + Co(s)
The reaction isA. spontaneous and Eo is positive.B. spontaneous and Eo is negative.C. non-spontaneous and Eo is positive.D. non-spontaneous and Eo is negative
13. Referring to the data booklet, which of the following can act as an oxidizing agent but not as a reducing agent?A. ZnB. Cl-
37
C. Sn2+
D. Fe3+
14. Which equation represents a redox reaction?A. Pb2+ + 2Cl- → PbCl2
B. CaO + CO2 → CaCO3
C. Mg + 2HCl → MgCl2 + H2
D. HCl + NaOH → NaCl + H2O 15. In a redox reaction, ClO- was converted to Cl- in a basic solution. The balanced half-reaction for this
process is
A. ClO- + H2O + 2e- → Cl- + 2OH-
B. ClO- + 2OH- → Cl- + 2e- + H2OC. ClO- + H2O → Cl- + 2e- + 2OH-
D. ClO- + 2OH- + 2e- → Cl- + H2O
16. Which of the following gives the correct oxidation numbers for the nitrogen atoms in all three chemical species?
N2 Li3N NO2-
A. 0 -3 +3B. -3 -3 +3C. 0 +3 -3D. -3 -3 -3
17. The oxidation number of manganese changes as MnO4- is converted to MnO2. How many electrons are gained or lost by the manganese during the change?A. 1e- lostB. 1e- gainedC. 3e- lostD. 3e- gained
18. Which of the solutions listed below should NOT be stored in a tin-plated container?I NaNO3(aq) II AgNO3(aq) III SnBr2(aq) IV Cl2(aq)
A. I onlyB. II and IVC. II, III and IVD. II and III
19. If fluorine gas is bubbled through NaI(aq),
A. Na+(aq) is reduced
B. I-(aq) is oxidized
C. F2(aq) is oxidizedD. I-
(aq) is reduced
20. A high school laboratory's waste container is used to dispose of aqueous solutions of sodium nitrate, potassium sulfate, hydrochloric acid, and tin (II) chloride. The most likely net redox reaction predicted to occur inside the waste container is represented by the equation:
A. Sn2+(aq) + 2NO3
-(aq) + 4H+
(aq) ® 2NO(g) + 2H2O(l) + Sn4+(aq)
B. 2 H+(aq) + 2K+
(aq) ® H2(g) + K(s)
C. SO42-
(aq) + 4H+(aq) + 2 Cl-
(aq) ® H2SO3(aq) + H2O(l) + Cl2(g)
D. Cl2(g) + Sn2+(aq) ® Cl-
(aq) + Sn(s)
38
21. The net ionic equation for the reaction between aqueous solutions of tin (II) bromide and iron (III) nitrate is
A. Sn2+(aq) + NO3
-(aq) ® NO2(g) + Sn4+
(aq)
B. Sn2+(aq) + 2 Fe3+
(aq) ® Sn4+(aq) + 2 Fe2+
(aq)
C. 2Sn2+(aq) + Fe3+
(aq) ® 2 Sn3+(aq) + Fe2+
(aq)
D. Sn2+(aq) + 2NO3
-(aq) ® 2NO2(g) + Sn4+
(aq
22. Use the following unbalanced redox reaction to answer this question.
__O2(g) + ___H+(aq) + ___Ag(s) ® __Ag+
(aq) + ___H2O(l)
The coefficients for the balanced equation areA. 1,4,4,4,2B. 1,4,1,1,2C. 1,2,1,1,1D. 1,1,1,1,1
23. A student mixes various combinations of the metals Pd, Cd and Ga with 1.0 M solutions of their ions. The student then measures the voltages and puts them in the table below.1 Ga Pd2+ +1.18 V2 Ga Cd2+ +0.16 V3 Cd Pd2+
What voltage is produced by reaction # 3?A. -1.34 VB. -1.02 VC. +1.02 VD. +1.34 V
24. Which of the following is a correctly balanced reduction half-reaction?A. 2HCN + 2e- ® C2N2 + 2H+
B. 2Sb + 3H2O + 6e- ® Sb2O3 + 6H+
C. NO3- + 3H+ + 3e- ® HNO2 + H2O
D. Sb2O5 + 6H+ + 4e- ® 2Sb(OH)2+ + H2O
25. Which of the following combinations will react spontaneously under standard conditions?A. Ag + Br2
B. Ni + Co2+
C. Zn + Mg2+
D. Au + HNO3
26. Which of the following best describes what happens when lead solid is placed in a 1.0 M solution of Cu(NO3)2?A. Bubbles form on the lead.B. No changes are observed.C. Copper solid forms on the lead and the solution changes colour.D. The mass of lead solid increases and the solution does not change colour.
27. Consider the following unbalanced equation for a redox reaction in acidic solution: Br- + BrO3
- ® Br2
What is the equation for the balanced reduction half-reaction?
39
A. 2Br- ® Br2 + 2e-
B. 2Br- + 2e- ® Br2
C. 5e- + 6H+ + 2BrO3- ® Br2 + 3H2O
D. 10e- + 12H+ + 2BrO3- ® Br2 + 6H2O
28. What reaction will occur when a solution containing 1.0 M MgSO4 and 1.0 M CoCl2 is stored in a galvanized (Zn coated) bucket?A. Mg(s) + Cl2 ® MgCl2(s)
B. Co2+ + SO42- ® CoSO4(s)
C. Co2+ + Zn(s) ® Zn2+ + Co(s)
D. Mg2+ + Zn(s) ® Zn2+ + Mg(s)
29. When the skeletal equation I- ® IO3- is balanced in acidic solution, H2O, H+
and e- will appear. Which of the following are the correct balancing coefficients?H2O H+ e-
A. 3 3 2B. 3 6 6C. 3 6 5D. 6 12 10
30. Consider the following balanced redox equation in acidic solution:5H2O + 2CoCl2 + OCl- ® 2Co(OH)3 + 5Cl- + 4H+
Which of the following describes the amounts and locations of OH- and H2O if the equation is balanced in basic solution?A. 1H2O on the left and no OH-
B. 1H2O on the left and 4OH- on the leftC. 5H2O on the left and 4OH- on the leftD. 1H2O on the left and 4OH- on the right
Redox Quiz #4 Electrochemical Cells/Electrolytic Cells
1.0 M KNO3
1M Zn(NO3)2
PbZn
1M Pb(NO3)2
voltmeter
40
1. In the electrochemical call above, the electrons flow fromA. zinc to lead and the mass of zinc increasesB. zinc to lead and the mass of lead increasesC. lead to zinc and the mass of zinc increasesD. lead to zinc and the mass of lead increases
2. The initial cell voltage isA. -0.89 VB. -0.63 VC. +0.63 VD. +0.89 V
3. In an operating lead-zinc electrochemical cell shown above, the cathodeA. gains mass as anions are reducedB. loses mass as anions are reducedC. gains mass as cations are reducedD. loses mass as cations are reduced
4. The equation for the half-reaction at the anode isA. Zn2+ + 2e- → ZnB. Pb2+ + 2e- → PbC. Zn → Zn2+ + 2e-
D. Pb → Pb2+ + 2e-
5. The equation for the half-reaction at the cathode is
A. Zn2+ + 2e- → ZnB. Pb2+ + 2e- → PbC. Zn → Zn2+ + 2e-
D. Pb → Pb2+ + 2e-
6. The direction of electron flow in an electrochemical cell is fromA. anode to cathode through the external wireB. cathode to anode through the external wireC. anode to cathode through the external wire and back through the salt bridgeD. cathode to anode through the external wire and back through the salt bridge
7. Which of the following is formed at the anode during the electrolysis of 1.0 M NaI?A. I2
B. O2
C. H2
D. Na
Power Source
- +
PtPt
41
Molten MgCl2
8. As this cell operatesA. Cl- is oxidized at the anodeB. Mg2+ is oxidized at the anodeC. Cl- is oxidized at the cathodeD. Mg2+ is oxidized at the cathode
9. In an operating electrochemical cell, the anions migrateA. towards the anode through the wireB. towards the cathode through the wireC. towards the anode through the salt bridgeD. towards the cathode through the salt bridge
10. As the above electrochemical cell operatesA. nitrate ions migrate into the copper half-cellB. copper (II) ions migrate through the salt bridgeC. magnesium ions migrate through the salt bridgeD. potassium ions migrate into the magnesium half-cell
11. In the above electrochemical cell, the reaction at the anode is A. Cu → Cu2+ + 2e-
B. Cu2+ + 2e- → CuC. Mg → Mg2+ + 2e-
D. Mg2+ + 2e- → Mg
12. In the above electrochemical cell, the initial voltage isA. 2.03 VB. 2.52 VC. 2.71 VD. 2.89 V
13. Which of the following aqueous solutions produces H2(g) and O2(g) during electrolysisA. 1.0 M KIB. 1.0 M CuI2
C. 1.0 M K2SO4
D. 1.0 M CuSO4
1.0 M KNO3
1 M Mg(NO3)2
CuMg
1 M CuSO4
voltmeter
42
14. In the electrolysis of molten zinc chloride, the half-reaction at the anode isA. Cl2 + 2e- → 2Cl-
B. 2Cl- → Cl2 + 2e-
C. Zn2+ 2e- → ZnD. Zn → Zn2+ + 2e-
15. The initial cell voltage at 25oC isA. -1.06 VB. -0.54 VC. +0.54 VD. +1.06 V
16. The balanced equation for the overall reaction isA. Ni+
(aq) + Ag(s) → Ag+(aq) + Ni(s)
B. Ni(s) + Ag+(aq) → Ag(s) + Ni+
(aq)
C. Ni2+(aq) + 2Ag(s) → 2Ag+
(aq) + Ni(s)
D. Ni(s) + 2Ag+(aq) → 2Ag(s) + Ni2+
(aq)
17. This redox reaction occurs becauseA. Ag(s) is a stronger oxidizing agent than Ni(s)
B. Ag(s) is a weaker reducing agent than Ni(s)
C. Ag+(aq) is a stronger reducing agent than Ni2
+(aq)
D. Ag+(aq) is a weaker oxidizing agent than Ni2
+(aq)
1.0 M KNO3
1 M Ni(NO3)2
AgNi
1 M AgNO3
voltmeter
1.0 M KNO3
1 M Au(NO3)3
PbAu
1 M Pb(NO3)2
voltmeter
43
18. The direction of the electron flow isA. from Au to Pb through the wireB. from Pb to Au from the wireC. from Au to Pb through the salt bridgeD. from Pb to Au through the salt bridge
19. As the cell operatesA. NO3
- and K+ will migrate toward the Pb half-cellB. NO3
- and K+ will migrate toward the Au half-cellC. NO3
- will migrate toward the Pb half-cell and K+ will migrate toward the Au half-cellD. NO3
- will migrate toward the Au half-cell and K+ will migrate toward the Pb half-cell
20. The initial voltage isA. -1.37 VB. 0.00 VC. 1.37 VD. 1.63 V
21. Which of the following is a balanced half-reaction in base?A. Cl2 + 3H2O → ClO3
- + 6H+ + 5e-
B. Cl2 + 6OH- → ClO3- + 5e- + 3H2O
C. Cl2 + 6H2O → 2ClO3- + 12H+ + 10e-
D. Cl2 + 12OH- → 2ClO3- + 6H2O + 10e-
22. In which of the following unbalanced equations does chromium undergo oxidation?A. Cr3+ → CrB. Cr3+ → Cr2+
C. Cr3+ → Cr2O72-
D. CrO42- → Cr2O7
2-
23. Which of the following is formed at the anode and cathode during the electrolysis of 1.0 M Na2SO4?
Anode CathodeA. O2 H2
B. H2 O2
C. H2 NaD. S Na
24. Consider the electrolytic cell shown in the following diagram:
Power Source
- +
CuPt
1 M AgNO3
44
Which of the following describes the anion movement and electrode masses for the above cell?
Anion Movement Mass of Pt Electrode Mass of Cu ElectrodeA. to the Cu increases increasesB. to the Cu increases decreasesC. to the Pt decreases increasesD. to the Pt decreases decreases
25. Which of the following occurs when 1 M NiSO4 is electrolyzed using inert carbon electrodes?A. The cathode dissolves.B. Hydrogen gas is produced.C. The pH of the solution decreases.D. The [Ni2+] increases.
26. Which of the following describes electrolysis?
A. a process that uses electrical energy to cause a spontaneous reactionB. a process that generates electrical energy using a spontaneous reactionC. a process that uses electrical energy to cause a non-spontaneous reactionD. a process that generates electrical energy using a non-spontaneous reaction
Power Source
- +
CuPt
1 M NiI2
45
27. What products would form at the anode and cathode as this cell operates?Anode Cathode
A. I2 NiB. Ni I2
C. O2 H2
D. Cu2+ Ni
28. In the above cell, if 1.0 M NiI2 is replaced with molten NiI2, what products would form at the electrodes?Anode Cathode
A. I2 NiB. Ni I2
C. O2 H2
D. Cu2+ Ni
29. In the above cell, which describes the movement of the electrons?A. They move from left to right towards the anode.B. They move from right to left towards the anode.C. They move from right to left towards the cathode.D. They move from left to right towards the cathode.
30. Which of the following best describes what happens to the mass of the anode and the mass of the cathode as the cell operates?
Anode Mass Cathode MassA. decreases increasesB. decreases no changeC. no change decreasesD. no change no change
31. What is the standard voltage E0 for the cell?A. 0.43 VB. 0.77 V
1.0 M KNO3
1 M KMnO4Acidified
PtPt
1 M H2SO3
voltmeter
46
C. 1.34 VD. 1.68 V
Quiz #5 Application of Cells 1. The corrosion of iron can be prevented by attaching a piece of zinc to the iron because
theA. iron acts as an anodeB. zinc reduces more readily than ironC. electrons flow from the zinc to the ironD. iron ions form more readily than zinc ions
2. An iron spoon is electroplated with copper. The equation representing the reduction reaction is
A. Cu2+(aq) + 2e- → Cu(s)
B. Cu(s) → Cu2+(aq) + 2e-
C. Fe2+(aq) + 2e- → Fe(s)
D. Fe(s) → Fe2+(aq) + 2e-
3. In an operating zinc-copper electrochemical cell, the oxidizing agent
A. loses electrons at the anodeB. loses electrons to the cationsC. gains electrons at the cathodeD. gains electrons from the anions
4. An example of electro refining is the
A. extraction of aluminum from bauxiteB. purification of lead from an impure anodeC. recovery of zinc from a zinc sulphide solutionD. production of chlorine from a sodium chloride solution
5. Electroplating always involves the
A. oxidation of anionsB. reduction of cationsC. reduction at the anodeD. oxidation at the cathode
6. Hydrogen and oxygen react to provide energy in a
A. dry cellB. fuel cellC. alkaline cellD. lead-acid storage cell
7. En electrolytic process is used to purify impure lead. The electrodes are
ANODE CATHODEA. carbon impure leadB. pure lead carbonC. pure lead impure leadD. impure lead pure lead
47
8. In the cell below the half-reaction at the cathode is
A. Cu2+ + 2e- → Cu(s)
B. 2SO42- → S2O8
2- + 2e-
C. H2O → ½ O2(g) + 2H+ + 2e-
D. 2H2O + 2e- → H2(g) + 2OH-
9. In the electrolysis of molten PbBr2, the products at the anode and cathode are
ANODE (INERT)
CATHODE (INERT)
Br2 H2
O2 PbPb Br2
Br2 Pb
10. Under which conditions could an electrochemical cell provide 0.93V?
A.B.C.D.
Power Source
- +
Pt
Iron Key
1.0 M CuSO4
48
Anode Cathode
Cu MgMg CuAg PbPb Ag
11. The reduction reaction in the above electrochemical cell is
A. Pb2+ + 2e- → PbB. Pb → Pb2+ + 2e-
C. Ag+ + e- → AgD. Ag → Ag+ + e-
12. An industrial process involving electrolysis is the reduction of
A. water forming oxygen gasB. water forming hydrogen gasC. sea water forming chlorine gasD. sea water forming bromine liquid
13. To plate a nickel coin with copper
A. the nickel coin must be the cathodeB. the cathode must be made of copperC. the electrons must flow to the anodeD. the solution must contain nickel ions
14. Which of the following ions can be reduced from an aqueous solution
A. Ba2+
B. Al3+
C. Sn2+
D. Na+
15. The principal function of a fuel cell is to
A. produce fuelB. electrolyze fuelC. produce hydrogenD. produce electricity
16. If a piece of nickel is to be gold-plated using an electrolytic process, which half-reaction occurs at the
cathode?
A. Ni → Ni2+ + 2e-
B. Ni2+ + 2e- → NiC. Au → Au3+ + 3e-
D. Au3+ + 3e- → Au
A.B.C.D.
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17. Consider the following redox reactionAs2O3 + 2NO3
- + 2H2O + 2H+ → 2H3AsO4 + N2O3
In this reaction, nitrogenA. loses electrons and increases in oxidation numberB. gains electrons and increases in oxidation numberC. loses electrons and decreases in oxidation numberD. gains electrons and decreases in oxidation number
18. In an electrochemical cell, the cathodeA. is reducedB. loses massC. is the reducing agentD. is the site of reduction
19. 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 isA. 2I- → I2 + 2e-
B. Na+ + e- → NaC. H2O + ½ O2 + 2H+ + 2e-
D. 2H2O +2e- → H2 + 2OH-
Redox Quiz #6 Corrosion & Cathodic Protection Titration
1. Which of the following metals could be used to cathodically protect a sample of lead?A. ironB. goldC. silverD. copper
2. A piece of iron can be prevented from corroding byA. making it a cathodeB. placing it in an acidic solutionC. attaching a small piece of lead to itD. attaching a small piece of gold to it
3. To determine the [Fe2+] in a solution of FeSO4 by e redox titration, a suitable reagent would be an acidified solution ofA. Cr3+
B. Mn2+
C. SO42-
D. Cr2O72-
4. As a metal corrodes,A. it gains electronsB. it becomes reducedC. it acts as a reducing agentD. its oxidation number decreases
5. Which method will cathodically protect a piece of iron?A. Paint the iron
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B. Cover the iron with greaseC. Attach a piece of lead tot he ironD. Attach a piece of magnesium to the iron
6. Corrosion of iron can be prevented by attaching a piece ofA. MnB. CuC. PbD. Sn
7. 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 isA. less than 0.34 VB. greater than 1.50 VC. greater than 0.85 V but less than 1.50 VD. greater than 0.34 V but less than 0.85 V
8. Consider the following redox equation:Br2 + SO2 + Na2SO4 + 2H2O → 2H2SO4 + 2NaBr
Which of the following is gaining electrons?A. Br2
B. SO2
C. H2OD. Na2SO4
9. The reaction that occurs when pieces of lead, zinc, copper and silver are placed in a solution of Ni(NO3)2 isA. Pb + Ni2+ → Pb2+ + NiB. Zn + Ni2+ → Zn2+ + NiC. Cu + Ni2+ → Cu2+ + NiD. 2Ag + Ni2+ → 2Ag+ + Ni
10. In the electrochemical cell above, the electrons flow fromA. copper to lead through the wireB. lead to copper through the wireC. copper to lead through the salt bridgeD. lead to copper through the salt bridge
11. In the electrochemical cell above, the initial Eo value is
1.0 M KNO3
1M Pb(NO3)2
CuPb
1M Cu(NO3)2
voltmeter
51
A. 0.03 VB. 0.21 VC. 0.29 VD. 0.47 V
12. A reaction that occurs during the corrosion of iron isA. Fe + 3e- → Fe3+
B. Fe → Fe2+ + 2e-
C. Fe2+ + 2e- → FeD. Fe3+ + e- → Fe2+
13. Consider the following reactionZn(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. 0.0131 LB. 0.0262 LC. 0.200 LD. 0.400 L
14. Consider the following diagram:
Why would this cell fail to electroplate the Fe nail with copper?A. The Cu is inert.B. The Fe nail is the anode.C. The Fe nail is the cathode.D. The porous membrane prevents reaction.
Power Source
- +
FeCu
1 M Cu(NO3)2
Porous Membrane
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15. Consider the following diagram:
Why would this cell fail to electroplate the Fe nail with copper?A. The Cu is the cathode.B. The electrolyte does not contain Cu2+.C. The Fe nail is the cathode.D. The porous membrane prevents reaction.
16. A 10.0 mL water sample was analyzed for [Fe2+] using a redox titration with acidified KMnO4. The equation for the reaction is:MnO4
- + 5Fe2+ + +8H+ ® Mn2+ + 5Fe3+ + 4H2O A 10.0 mL sample was titrated with 12.5 mL of 0.100 M KMnO4 solution.
What is the [Fe2+] in the water sample?A. 0.025 MB. 0.13 MC. 0.28 MD 0.63 M
17. Why is aluminum a good choice for the manufacture of outdoor structures? A. Pure aluminum is easily reduced.
B. Pure aluminum is not easily oxidized.C. Pure aluminum is easily reduced, but forms a protective coating.D. Pure aluminum is easily oxidized, but forms a protective coating.
18. Which of the following are produced at the anode and the cathode during the electrolysis of aqueous calcium iodide using carbon electrodes?
Anode CathodeA. Iodine CalciumB. Hydrogen OxygenC. Oxygen HydrogenD. Iodine Hydrogen
19. Which of the following are produced at the anode and the cathode during the electrolysis of molten calcium iodide using carbon electrodes?
Anode CathodeA. Iodine CalciumB. Hydrogen OxygenC. Oxygen HydrogenD. Iodine Hydrogen
Power Source
+ -
FeCu
1 M AgNO3
Porous Membrane
53
20. Which of the following are produced at the anode and the cathode during the electrolysis of aqueous potassium fluoride using carbon electrodes?
Anode CathodeA. Oxygen PotassiumB. Hydrogen OxygenC. Oxygen HydrogenD. Fluorine Potassium
21. Which of the following are produced at the anode and the cathode during the electrolysis of molten potassium fluoride using carbon electrodes?
Anode CathodeA. Oxygen PotassiumB. Hydrogen OxygenC. Oxygen HydrogenD. Fluorine Potassium
22. Two reactions involved in the refining of copper are: Reaction I 2Cu2S + 3O2 ® 2Cu2O + 2SO2
Reaction II Cu2S + 2Cu2O ® 6Cu + SO2
What happens to the copper ions in this process?A. They are reduced in Reaction I.B. They are reduced in Reaction II.C. They are oxidized in Reaction I.D. They are oxidized in Reaction II.
23. Identify the oxidation number for manganese in MnO4-.
A. -7B. 7C. 8D. 9
24. Which of the following is more difficult to reduce than the H+ ion? A. I2
B. Ag+
C. Zn2+
D. Cu2+
25. The equation for the decomposition of nitrous acid is3HNO2 ® 2NO + HNO3 + H2O Which of the following is correct?A. This is a redox reaction.B. This is an acid-base reaction.C. This is a reduction half equation.D. This is an oxidation half equation.
26. An equation for the rusting of iron is shown below:4Fe + 3O2 ® 2Fe2O3
Which of the following is false?A. This is a redox reaction.B. O2 is the oxidizing agent.C. Metallic iron is reduced to Fe3+.D. Metallic iron is the reducing agent.
54
27. In which of the following chemical changes will there be an oxidation number change of +3 ?A. Cr3+ ® Cr2+
B. ClO- ® ClO2-
C. Cr3+ ® Cr2O72-
D. Mn2+ ® MnO4-
28. Consider the following spontaneous reactions:Cd2+ + Np ® Cd + Np3+
Cd + Pd2+ ® Pd + Cd2+
Np3+ + Ce ® Np + Ce3+
Which is the strongest oxidizing agent?A. Cd2+
B. Ce3+
C. Np3+ D. Pd2+
29. Consider the following equation:H3AsO4 + 4Zn + 8H+ ®AsH3 + 4Zn2+ + 4H2OWhich of the following is correct?A. Oxygen is reduced.B. Arsenic is reduced.C. Zinc is the oxidizing agent.D. The reaction is not a redox reaction.
30. What is the oxidation number of iron in magnetite, Fe3O4 ?A. +4/3
B. +2C. +8/3
D. +3
Redox Web Review
1) Which most readily gains electrons?
Cu Cu2+ Fe2+2 Zn2+ Au3+
2) 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
55
4. 2MnO4- + 5Pb +16H+ → 2Mn2+ + 8H2O + 5Pb2+
5. Will AgNO3 react with Zn? Write a balanced redox reaction and calculate Eo
6. 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 Electrolytic cell
8. What are the similarities between an electrochemical cell and an electrolytic cell?
Electrochemical cell or Electrolytic cell
56
9. State how you would determine each of the following in an electrochemical or electrolytic cell.
Electrochemical Cell Electrolytic 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
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?
57
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 → 3V + 2Al3+ Eo = +0.46 V
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) → HV2O73-
58
16. Balance the following redox reaction occurring in basic solution:
MnO4- + C2O4
2- → MnO2 + CO2
17. 250.0 ml 0.200M MnO4- reacts with excess SO32-. How many grams of MnO2 are
produced?
2MnO4- + 3SO3
2- +H2O → 2MnO2 +3SO42- + 2OH-
18. Determine the oxidation number for each bold atom.
MnO2 IO3- Cr2O7
2- C2O42- Al(NO3)3 NH4Cl NaH
HOOH NO3- H3PO4 Na2C2O4 I2 N2O3 Pt(H2O)4
2+
19. 250.0 mL of 0.500M MnO4- are required to titrate a 100.0 ml sample of SO3
-2. Calculate the [SO3
-2]
2MnO4- + 3SO3
2- +H2O → 2MnO2 + 3SO42-+ 2OH-
59
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
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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+ + Cu →
26) 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. Which of the above cells requires continuous input of O2 and H2 and is produced by Ballard Industries.
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28. 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
29. Describe each term:
salt bridge
electrolyte
anode
cathode
spontaneous
electron affinity
cation
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anion
electrochemical cell
electrolytic cell
oxidation number
electrolysis
oxidation
reduction
oxidizing agent
reducing agent
electrode
corrosion
electrowinning
electrorefining
over potential effect
fuel cell
30. 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.
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31. Which you would choose Zn or Cu to cathodically protect iron?
32. 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?
33. Completely analyze the following electrochemical cell.
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:
1.0 M KNO3
1 M Zn(NO3)2
CuZn
1 M Cu(NO3)2
voltmeter
64
Describe the change in [Cu+2] in the Cu half cell
Describe the change in [NO3-1] in the Zn half cell
34. Completely analyze the following electrochemical cell.
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
1 M HCl
Cu
1 M Cu(NO3)2
voltmeter
H2(g)
1.0 M KNO3
Pt
65
35. Completely analyze the following electrolytic cell.
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 ?
Power Source
- +
C
C
Molten Al2O3
66
36. Completely analyze the following electrolytic cell. Note that the electrodes are not inert and because of that, the anode might oxidize.
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 ?
Power Source
- +
Cu
Cu
1 M NaF
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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. FeB. COC. CO2
D. Fe2O3
2. Consider the following reaction: 2HNO2 + 2I- + 2H+ → 2NO + I2 +2H2O
The oxidation number for each nitrogen atomA. increases by 1B. increases by 2C. decreases by 1D. decreases by 2
3. Which of the following reactions is spontaneous?
A. 2I- + Ag → Ag+ + I2
B. Co2+ + Cu → Co + Cu2+
C. Cu2+ + Pb → Pb2+ + CuD. Ni2+ + 2Ag → 2Ag+ + Ni
4. Consider the following redox reaction for a lead-acid storage cell:
Pb + PbO2 + 4H+ + 2SO42- → 2PbSO4 + 2H2O
The balanced, reduction half reaction isA. Pb + SO4
2- → 2PbSO4 + 2e-
B. Pb + 2H+ + SO42- → PbSO4 + 2H2O + 2e-
C. PbO2 + 4H+ + SO42- + 2e- → PbSO4 + 2H2O
D. PbO2 + 2SO42 + 2H2O + 2e- → PbSO4 + 2OH-
5. 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. -1.12 VB. -0.40 VC. +0.40 VD. +1.12 V
6. Which of the following involves a nonspontaneous redox reaction?
A. fuel cellB. electroplatingC. redox titrationD. carbon dry cell
7. 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 molesB. 0.060 molesC. 1.5 moles
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D. 0.30 moles
8. Which of the following is not a redox reaction?A. Cu + Br2 → CuBr2
B. CO + H2O → CO2 + H2
C. CH4 + H2O → CO2 + 2H2OD. NaOH + HCl → NaCl + H2O
9. What is the minimum voltage required to form nickel from an aqueous solution of NiI2 using inert
electrodes?A. 0.26 VB. 0.28 VC. 0.54 VD. 0.80 V
10. What substances are formed at the anode and cathode during electrolysis of molten sodium chloride?
Anode CathodeA. O2 H2
B. Na Cl2
C. Cl2 H2
D. Cl2 Na 11. A solution containing an unknown cation reacts spontaneously with both zinc and copper. The unknown
cation isA. 1.0 M H+
B. 1.0 M Ag+
C. 1.0 M Sr2+
D. 1.0 M Mn2+
12. Which of the following half-reactions are balanced?
A. ClO- + H2O + e- → Cl2 + 2OH-
B. 2ClO- + H2O + 2e- → Cl2 + 3OH-
C. 2ClO- + 2H2O + 2e- → Cl2 + 4OH-
D. 2ClO- + 2H2O → Cl2 + 4OH- + 2e- 13. Which of the following is a spontaneous redox reaction?
A. Ag+ + I- → AgIB. Ag+ + Fe2+ → Ag + Fe3+
C. 3Ag+ + Au → 3Ag + Au3+
D. 2Ag+ + Ni2+ → 2Ag + Ni 14. Salting the roads during the winter increases the amount of corrosion of cars. The is because the salt
A. reacts with the ironB. provides an electrolyteC. acts as a reducing agent D. acts as an oxidizing agent
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Consider the following electrochemical cell for the next five questions. 15. The half-reaction that occurs at the anode is
A. Ni → N2+ + 2e-
B. Ni2+ + 2e- → NiC. Cu → Cu2+ + 2e-
D. Cu2+ + 2e- → Cu 16. The half-reaction that occurs at the cathode is
A. Ni → N2+ + 2e-
B. Ni2+ + 2e- → NiC. Cu → Cu2+ + 2e-
D. Cu2+ + 2e- → Cu 17. The cell potential or Eo is
A. 0.41 VB. 0.78 VC. 0.34 VD. 0.60 V
18. The following ions migrate to the Cu electrode
A. K+ Cu2+ Ni2+
B. Cu2+ Ni2+
C. Cl- NO3-
D. Cl- NO3- 2e-
19. The electrons flow
A. through the salt bridge from Cu to NiB. through the salt bridge from Cu to NiC. through the wire from Cu to NiD. through the wire from Ni to Cu
20. Which of the following will not react spontaneously with 1.0 M HCl?
A. tinB. lithiumC. mercuryD. magnesium
NiCuu
1.0 M KCl
1M Cu(NO3)2 1M Ni(NO3)2
voltmeter
70
21. Which of the following can be produced by electrolysis from a 1.0 M aqueous solution containing its ion?A. nickelB. sodiumC. aluminumD. magnesium
22. In order for an electrolytic cell to operate, it must have
A. a voltmeter.B. a salt bridge.C. a power supply.D. an aqueous solution.
23. In the electrolysis of molten ZnCl2 using carbon electrodes, the reaction that occurs at the anode is
A. Zn → Zn2+ + 2e-
B. Zn2+ + 2e- → ZnC. 2Cl- → Cl2 + 2e-
D. Cl2 + 2e- → 2Cl-
24. In the electrolysis of molten zinc chloride, the half-reaction at the anode is
A. Cl2 + 2e- → 2Cl-
B. 2Cl- → Cl2 + 2e-
C. Zn2+ + 2e- → ZnD. Zn → Zn2+ + 2e-
25. The corrosion of iron can be prevented by attaching a piece ofA. MnB. CuC. PbD. Sn
26. The oxidation number of carbon in CaC2O4 is
A. +2B. +3C. +4D. +6
27. To plate a nickel coin with copper,
A. the nickel coin must be the cathode.B. the cathode must be made out of copperC. the electrons must flow to the anodeD. the solution must contain nickel ions
Consider the following electrochemical cell for the next five questions.
ZnCuuu
1.0 M KNO3
1M Cu(NO3)2 1M Zn(NO3)2
voltmeter
71
28. Which of the following statements apply to this electrochemical cell?I Electrons flow through the wire toward the copper electrode.II The copper electrode increases in mass.III Anions move toward the Zn half-cell.A. I and II onlyB. I and III onlyC. II and III onlyD. I, II, and III
29. The balanced equation for the overall reaction isA. Zn + Cu2+ → Cu + Zn2+
B. Cu + Zn2+ → Zn + Cu2+
C. Zn2+ + Cu → Cu2+ + ZnD. Cu + Zn → Zn + Cu
30. At equilibrium the voltage of the above cell is
A. -1.10 VB. 0.00 VC. +0.42 VD. +1.10 V
31. This redox reaction occurs because
A. Zn is a stronger oxidizing agent than CuB. Zn is a stronger reducing agent than Cu,C. Cu is a stronger oxidizing agent than ZnD. Zn2+ is a weaker reducing agent than Cu2+
32. The initial cell voltage at 25 oC is
A. -1.10 VB. +1.10 VC. +0.91 VD. +0.86 V
33. Consider the following redox reaction: Co2+
(aq) + 2Ag(s) ⇋ 2Ag+(aq) + Co(s)
The reaction isA. spontaneous and Eo is positiveB. spontaneous and Eo is negativeC. non-spontaneous and Eo is positiveD. non-spontaneous and Eo is negative
34. When MnO4
- reacts to form Mn2+, the manganese in MnO4- is
A. reduced as its oxidation number increasesB. reduced as its oxidation number decreasesC. oxidized as its oxidation number increasesD. oxidized as its oxidation number decreases
35. The electrolyte used in the alkaline battery is
A. KClB. NaOHC. H2SO4
D. KOH
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36. The electrolyte used in an automobile battery isA. KClB. NaOHC. H2SO4
D. KOH 37. The anode used in the commercial production of Aluminum is
A. CB. PtC. AlD. Al2O3
38. The anode and cathode used in the electrorefining of impure lead to pure lead are
Anode CathodeA. Pure Pb Impure PbB. Impure Pb Pure PbC. Pb2+ PbD. Pb Pb2+
39. The anode in the LeClanche or common dry cell is
A. CB. ZnC. MgD. KOH
40. Which of the following are electrolytic cells
I Electro winningII ElectroplatingIII Charging a car batteryIV Fuel cell
A. I and II onlyB. I, II, and III onlyC. II and II onlyD. I, II, III, and IV
Subjective 1. Balance the following in basic solution.
MnO4- + C2O4
2- → MnO2 + CO2 (basic)
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2. 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. 3. Consider the following diagram for the electro refining of lead. 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:
Power Source
Impure PbPure Pb
74
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 2.00 13.80 24.55Final burette reading 13.80 24.55 35.32
6. 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 electrodes are not inert
and can oxidize: Cl-, or Cu will oxidize)
a) Identify the product at the anode
b) Write the equation for the reduction half-reaction.
c) Write the equation for the overall reaction. 8. Completely analyze the following electrochemical cell.
MnSnnnn
1.0 M KNO3
1M Sn(NO3)2 1M MnNO3)2
voltmeter
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Chemistry 12 Electrochemistry Practice Test 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 Pd → Pd2+ + 2e-
B Ni → Ni2+ + 2e-
C Pd2+ + 2e- → PbD Ni2+ + 2e- → Ni
2. As the cell operates,
A both the K+ and the NO3- migrate into the nickel half-cell
B both the K+ and the NO3- migrate into the palladium half-cell
C the K+ migrates into the nickel half-cell and the NO3- migrates into the palladium half-cell
D the K+ migrates into the palladium half-cell and the NO3- migrates into the nickel half-cell
3. The initial cell voltage is 1.21 V. The reduction potential of Pd2+ is
A -1.21 VB -.95 VC +0.95D +1.21 V
4. What substances are formed at the anode and cathode during electrolysis of molten sodium chloride,
NaCl(l)?
Anode Cathode
A O2 H2
B Na Cl2
C Cl2 H2
D Cl2 Na
NiPd
1.0 M KCl
1M Pd(NO3)2 1M Ni(NO3)2
voltmeter
76
5. Consider the following electrolytic cell:
In the cell above
A I- migrates to the anode and gains electronsB I- migrates to the cathode and loses electronsC Na+ migrates to the anode and loses electronsD Na+ migrates to the cathode and gains electrons
6. Which of the following are necessary for electroplating to occur using an electrolytic cell?
I Two electrodesII A metal being reducedIII A direct current power supply
A I and II onlyB I and III onlyC II and III onlyD I, II, and III
7. A fuel cell consumes H2 and O2 gas, uses a KOH electrolyte, and produces electricity. The reaction at the
anode isA 2H+ + 2e- → H2
B 1/2O2 + 2H+ + 2e- → H2OC 4OH- → O2 + 2H2O + 4e-
D H2 + 2OH- → 2H2O + 2e-
8. A student investigating redox reactions recorded the following results:
V2+ + Te2- → no reactionU4+ + Te2- → U3+ + Te
Based on these results, the strengths of the oxidizing agents, arranged from strongest to weakest, are
A V2+ Te U4+
B U4+ Te V2+
C U3+ Te2- V2+
D V2+ Te2- U3+
9. What is the minimum voltage required to form nickel from 1 M NiI2 using inert electrodes?
Power Source+ -
Inert Electrode
Inert Electrode
Molten NaI(l)
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A 0.26 VB 0.28 VC 0.54 VD 0.80 V
10.
Which of the following occurs as the cell operates?A the Zn electrode is reduced and increases in massB the Zn electrode is reduced and decreases in massC the Zn electrode is oxidized and increases in massD the Zn electrode is oxidized and decreases in mass
11. Which of the following reactants would produce an E0 of +0.63 V?
A Ag+ + I2
B Pb2+ + ZnC Mg2+ + CaD Zn2+ + Mn
12. The concentration of Fe2+
(aq) can be determined by a redox titration usingA KBrB SnCl2
C KMnO4 (basic)D KBrO3 (acidic)
13. Which of the following will oxidize Fe2+?
A I2(s)
B Ni(s)
C Zn(s)
D Br2(l)
14. The oxidation number of carbon in C2O42- is
A +3B +4C +5D +6
15. Consider the following reaction: 3As2O3 + 4NO3- + 7H2O → 6H3AsO4 + 4NO
The oxidizing agent isA H+
NiZn
1.0 M KNO3
1M Zn(NO3)2 1M Ni(NO3)2
voltmeter
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B H2OC NO3
-
D AsO3
16. When W2O5 is converted to WO2 in a redox reaction, the W has been
A reduced since its oxidation number has increasedB reduced since its oxidation number has decreasedC oxidized since its oxidation number has increasedD oxidized since its oxidation number has decreased
17. Consider the following:
I WaterII Oxygen gasIII NitrogenAt 25oC, a piece of iron rusts in the presence of A I onlyB III onlyC I and II onlyD II and III only
18. Which of the following represents a redox reaction?
A H2CO3 → H2O + CO2
B CuS + H2 → H2S + CuC AgNO3 + NaCl → AgCl + NaNO3
D 2HCl + Na2SO3 → 2NaCl + H2SO3
19. The following reaction occurs in an electrochemical cell:
3Cu2+ + Cr → 2Cr3+ + 3CuThe Eo for the cell isA 0.40 VB 0.75 VC 1.08 VD 2.50 V
20. During the corrosion of magnesium, the anode reaction is
A Mg → Mg2+ + 2e-
B Mg2+ + 2e- → MgC 4OH- → O2 + 2H2O + 4e-
D O2 + 2H2O + 4e- → 4OH-
21. A molten binary salt, ZnCl2, undergoes electrolysis. The cathode reaction is
A Zn → Zn2+ + 2e-
B 2Cl- → Cl2 + 2e-
C Cl2 + 2e- → 2Cl-
D Zn2+ + 2e- → Zn 22. Which of the following represents a redox reaction?
A CaCO3 → CaO + CO2
B SiCl4 + 2Mg → Si + 2MgCl2
C 2NaOH + H2SO4 → 2H2O + Na2SO4
D AgBr + 2S2O32- → Ag(S2O3)2
3- + Br-
23. The process of applying an electric current through a cell to produce a chemical change is called
A corrosionB ionizationC hydrolysis
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D electrolysis
80
24. A student investigating redox reactions recorded the following results:V2+ + Te2- → no reactionU4+ + Te2- → U3+ + TeBased on these results, the strengths of the oxidizing agents, arranged from strongest to weakest, areA V2+ Te U4+
B U4+ Te V2+
C U3+ Te2- V2+
D V2+ Te2- U3+
25. A spontaneous redox reaction occurs when Sn2+ is mixed with
A I2
B CuC H2SD Ag2S
26. Consider the redox reaction: 2BrO3
- + 10Cl- + 12H+ → Br2 + 5Cl2 + 6H2Othe oxidation half-reaction ivolved in this reaction isA 2Cl- → Cl2 + 2e-
B 2H+ → H2 + 2e-
C BrO3- + 6H+ + 5e- → ½ Br2 + 3H2O
D BrO3- + 6H+ → ½ Br2 + 3H2O + 5e-
27. Which of the following is not a redox reaction?
A Cu + Br2 → CuBr2
B CO + H2O → CO2 + H2
C CH4 + O2 → CO2 + 2H2OD NaOH + HCl → NaCl + H2O
28. During the electrolysis of 1.0 M Na2SO4, the reaction at the cathode is
A Na+ + 1e- → NaB 2SO4
2- → S2O82- + 2e-
C 2H2O → O2 + H+ + 4e-
D 2H2O + 2e- → H2 + 2OH-
29. An oxidizing agent will cause which of the following changes?
A PtO2 → PtOB PtO3 → PtO2
C Pt(OH)2 → PtD Pt(OH)2
2+ → PtO3
30. 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 Cd loses 2e- and forms Cd(OH)2(s)
B Cd gains 2e- and forms Cd(OH)2(s)
C NiO2 loses 2e- and forms Ni(OH)2(s)
D 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 nickelB sodiumC aluminumD magnesium
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32. In the electrolysis of molten ZnCl2 using carbon electrodes, the reaction that occurs at the anode isA Zn → Zn2+ + 2e-
B Zn2+ + 2e- → ZnC 2Cl- → Cl2 + 2e-
D Cl2 + 2e- → 2Cl-
33. In order for the electrolytic cell to operate, it must have
A a voltmeterB a salt bridgeC a power supplyD 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 no reaction
Pt2+ reaction reaction
b) List the oxidizing agents in order of strongest to weakest 2. 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 the formation of rust.
a)
b)
3. Consider the following redox reaction:
H2Se + SO42- + 2H+ → Se + H2SO3 + H2O
Calculate the Eo for the reaction.
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4. Balance the following redox reaction in basic solution:
Au + Cl- + O2 → AuCl4- + OH-
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
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.
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7. Consider the following diagram in the electro refining of lead: 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
Power Source
Impure Pb
Pure Pb
