Oxidation and Reduction Reactions Workbook
Oxidation and Reduction Reactions Workbook
Period/Topic Worksheets Quiz
1. Oxidation, Reduction, Agents, & Reactions. WS 1
2. Lab: The Strength of Oxidizing Agents.
3. Oxidation Numbers Spontaneous Reactions WS 2 1
4. Oxidation Numbers, Application to Reactions. WS 3
5. Balancing Redox Half Reactions Acid/Base. WS 4 2
6. Balancing Redox Reactions in Acid/Base. WS 5
7. Standard Potentials Using Chart. WS 6 3
8. Electrochemical Cells. WS 7
9. Electrochemical Cells Lab.
10. Electrolytic Cells. WS 8 4
11. Electrolytic Cells Lab.
12. Application of Electrochemical Cells
13. 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. Oxidation
2. Reduction
3. Oxidizing agent
4. 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. 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:
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 _____
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.
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
2. Pb + Fe3+ → Fe2+ + Pb2+
3. Cl2 + F- → F2 + 2Cl-
4. S2O82- + 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+
13. Determine the oxidation number for the element underlined.
PbSO4 __________ ClO3- __________
HP032- __________ 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. HPO32- __________
22. Na202 __________
23. CaH2 __________
24. NaIO3 __________
25. C4H12 __________
26. Al2(SO4)3 __________
27. Al3+ + Zn → Al + Zn2+
Substance oxidized __________ Oxidizing agent __________
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. S2O32- → SO42-
2. MnO4- → Mn2+
3. As → AsO43-
4. Cr3+ → Cr2O72-
5. Pb2+ → PbO2
6. SO42- → S
7. NO3- → NO
8. NO3- → NH4+
9. BrO3- → Br2
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+ ________________
2I- → I2 ________________
IO3- → I2 ________________
MnO4- → Mn2+ ________________
ClO2- → ClO- ________________
20. Cr2O72- + 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. 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- → O2 + Cr3+
9. TeO32- + 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 + Cr2O72-
14. Te + NO3- → TeO32- + N2O4
15. IO3- + Re → ReO4- + IO-
16. Pb2+ + IO3- → PbO2 + I2
17. Cr2O72- + 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. H2SO4 22. HSO4-
24. P4 23. NaH
25. UO3 24. Na2O2
26. U2O5 25. PbSO4
Worksheet # 6 Review
1. 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) Na
b) 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.
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 Spontaneous
Ni(NO3)2 & Fe Spontaneous
Cr(NO3)3 & Fe 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?
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.
SO42- + Br2 → S2O32- + BrO3-
18. Balance in basic solution MnO4- + H2S → S + MnO
19. Describe as spontaneous or non-spontaneous. Use your reduction potential chart.
a) ZnCl2 & Cu b) CuCl2 & NaCl
c) 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. 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. MnO4- (in acid) oxidizing agent 1.51 V
22. 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.
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
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
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 Cr2O72- IO3- C2O42- 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.
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- + 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:
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:
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.
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.
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?
Quiz #1 Agents, Spontaneous Reactions, Oxidation #’s, and Strength
1. In a redox reaction, the species that loses electrons
A. is oxidized
B. is called the cathode
C. gains mass at the electrode
D. decreases in oxidation number
2. Which of the following is the strongest oxidizing agent?
A. Cu2+
B. Pb2+
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, are
A. Ag, Pt, Au
B. Pt, Au, Ag
C. Au, Ag, Pt
D. Au, Pt, Ag
4. MnO4- + 5Fe2+ + 8H+ → Mn2+ + 5Fe3+ + 4H2O The oxidizing agent in the reaction is
A. Fe2+
B. Fe3+
C. Mn2+
D. MnO4-
5. MnO4- + 5Fe2+ + 8H+ → Mn2+ + 5Fe3+ + 4H2O
During the reaction, electrons transfer from
A. Fe3+ to Fe2+
B. Fe2+ to MnO4-
C. MnO4- to Fe2+
D. MnO4- to Mn2+
6. As an element is oxidized, its oxidation number
A. increases as electrons are lost
B. decreases as electrons are lost
C. increases as electrons are gained
D. decreases as electrons are gained
7. A solution of 1.0 M Pb(NO3)2 will not react with a container made of
A. Cu
B. Fe
C. Sn
D. Zn
8. A spontaneous redox reaction occurs when a piece of iron is placed in 1.0 M CuSO4. The reducing agent is
A. Fe
B. Cu2+
C. H2O
D. SO42-
9. A substance is oxidized when it
A. loses protons
B. gains protons
C. loses electrons
D. 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 that
A. Ti2+ is a weaker reducing agent than Sn2+
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 reaction
A. loses mass
B. is oxidized
C. is the oxidizing agent
D. 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 |Observation |
|1 |U3+ |Y |reaction |
|2 |V2+ |U |reaction |
|3 |V2+ |Y |reaction |
|4 |Y3+ |V |no reaction |
The oxidizing agents from the strongest to the weakest are
A. V2+, U3+, Y3+
B. U3+, V2+, Y3+
C. Y3+, U3+, V2+
D. 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. 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 nitrogen
A. increases by 2
B. increases by 4
C. increases by 8
D. does not change
3. Consider the following redox equation:
12H+(aq) + 2IO3-(aq) + 10Fe2+(aq) → 10Fe3+(aq) + I2(s) + 6H2O(l)
The reducing agent is
A. 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+ + H2O
B. VO2 + H2 → V2+ + H2O + le-
C. VO2 + 2H+ + le- → V2+ + H2O
D. 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 3
B. increases by 6
C. decreases by 3
D. decreases by 6
7. Consider the following unbalanced half-reaction HClO2 ⇄ HClO
The balanced half-reaction would have
A. 1 electron on the left
B. 1 electron on the right
C. 2 electrons on the left
D. 2 electrons on the right
8. The oxidation number of platinum in Pt(H2O)42+ is
A. +2
B. 0
C. +4
D. +1/2
9. Consider the following half-reaction: BrO- → Br- (basic)
The balanced equation for the half-reaction is
A. BrO- + 2H+ + 2e- → Br- + H2O
B. 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+ + 3H2O
The species that loses the electron is
A. H2O
B. MnO4-
C. CH3CHO
D. CH3COOH
11. Hydrogen has an oxidation number of –1 in
A. H2
B. NaH
C. H2O
D. KOH
12. Consider the following:
2NO3- + 4H+ + 2e- → N2O4 + 2H2O
This equation represents
A. reduction
B. oxidation
C. neutralization
D. decomposition
13. Which of the following half-reactions is balanced?
A. IO3- + 6H+ +5e- → I2 + 3H2O
B. IO3- + 6H+ + 4e- →1/2 I2 + 3H2O
C. 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 is
A. Al
B. MnO4-
C. Al(OH)4-
D. MnO2
15. Consider the following redox reaction:
6H+ + 6I- + ClO3- → 3I2 + 3H2O + Cl-
The reducing agent is
A. I-
B. I2
C. H+
D. ClO3-
16. Nitrogen has an oxidization number of zero in
A. N2
B. NO2
C. NH3
D. HNO3
17. When MnO4- reacts to form Mn2+, the manganese in MnO4- is
A. reduced as its oxidation number increases
B. reduced as its oxidation number decreases
C. oxidized as its oxidation number increases
D. oxidized as its oxidation number decreases
18. Consider the following reaction:
2HNO3 + 3H2S → 2NO + 3S + 4H2O
The nitrogen in HNO3 undergoes
A. reduction
B. oxidation
C. electrolysis
D. neutralization
19. The oxidation number in carbon in CaC2O4 is
A. +2
B. +3
C. +4
D. +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. Cl2
B. Cr3+
C. H2O
D. Cr2O72-
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 is
A. -0.86 V
B. -0.60 V
C. +0.60 V
D. +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 unchanged
B. The reduction potentials of all the half-cells would increase by 0.76 V
C. The reduction potentials of all the half-cells would have positive values D. 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 |METAL |
|1 |CuCl2 |Zn |
|2 |CuCl2 |Ag |
|3 |CuCl2 |Ni |
Reactions occur in
A. beaker 2 only
B. beakers 1, 2, and 3
C. beakers 1 and 2 only
D. beakers 1 and 3 only
4. Consider the following redox reaction:
3SO2 + 3H2O + ClO3- → 3SO42- + 6H+ + Cl-
The reduction half-reaction is
A. 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+
C. Mn2+ and Cr2+
D. MnO2 and CrO42-
6. Bromine, Br2, will react spontaneously with
A. 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 is
A. +0.06 V
B. +0.39 V
C. +1.21 V
D. +1.54 V
9. A solution of 1.0 M Co(NO3)2 should be stored in a container made of
A. tin
B. zinc
C. aluminum
D. magnesium
10. A strong oxidizing agent has a
A. weak attraction for electrons
B. strong attraction for electrons
C. weak ability to become reduced
D. strong ability to become oxidized
11. The two species which react spontaneously in acidic solutions are
A. 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 is
A. 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. Zn
B. Cl-
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- + H2O
C. ClO- + H2O → Cl- + 2e- + 2OH-
D. ClO- + 2OH- + 2e- → Cl- + H2O
Quiz #4 Electrochemical Cells/Electrolytic Cells
1. In the electrochemical call above, the electrons flow from
A. zinc to lead and the mass of zinc increases
B. zinc to lead and the mass of lead increases
C. lead to zinc and the mass of zinc increases
D. lead to zinc and the mass of lead increases
2. The initial cell voltage is
A. -0.89 V
B. -0.63 V
C. +0.63 V
D. +0.89 V
3. In an operating lead-zinc electrochemical cell shown above, the cathode
A. gains mass as anions are reduced
B. loses mass as anions are reduced
C. gains mass as cations are reduced
D. loses mass as cations are reduced
4. The equation for the half-reaction at the anode is
A. Zn2+ + 2e- → Zn
B. Pb2+ + 2e- → Pb
C. Zn → Zn2+ + 2e-
D. Pb → Pb2+ + 2e-
5. The equation for the half-reaction at the cathode is
A. Zn2+ + 2e- → Zn
B. Pb2+ + 2e- → Pb
C. Zn → Zn2+ + 2e-
D. Pb → Pb2+ + 2e-
6. The direction of electron flow in an electrochemical cell is from
A. anode to cathode through the external wire
B. cathode to anode through the external wire
C. anode to cathode through the external wire and back through the salt bridge
D. 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
8. As this cell operates
A. Cl- is oxidized at the anode
B. Mg2+ is oxidized at the anode
C. Cl- is oxidized at the cathode
D. Mg2+ is oxidized at the cathode
9. In an operating electrochemical cell, the anions migrate
A. towards the anode through the wire
B. towards the cathode through the wire
C. towards the anode through the salt bridge
D. towards the cathode through the salt bridge
| |
10. As the above electrochemical cell operates
A. nitrate ions migrate into the copper half-cell
B. copper(II) ions migrate through the salt bridge
C. magnesium ions migrate through the salt bridge
D. 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- → Cu
C. Mg → Mg2+ + 2e-
D. Mg2+ + 2e- → Mg
12. In the above electrochemical cell, the initial voltage is
A. 2.03 V
B. 2.52 V
C. 2.71 V
D. 2.89 V
13. 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
14. 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- → Zn
D. Zn → Zn2+ + 2e-
15. The initial cell voltage at 25oC is
A. -1.06 V
B. -0.54 V
C. +0.54 V
D. +1.06 V
16. The balanced equation for the overall reaction is
A. 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 because
A. 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)
18. The direction of the electron flow is
A. from Au to Pb through the wire
B. from Pb to Au from the wire
C. from Au to Pb through the salt bridge
D. from Pb to Au through the salt bridge
19. As the cell operates
A. NO3- and K+ will migrate toward the Pb half-cell
B. NO3- and K+ will migrate toward the Au half-cell
C. NO3- migrates toward the Pb half-cell and K+ will migrate toward the Au
D. NO3- migrates toward the Au half-cell and K+ will migrate toward the Pb
20. The initial voltage is
A. -1.37 V
B. 0.00 V
C. 1.37 V
D. 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+ → Cr
B. Cr3+ → Cr2+
C. Cr3+ → Cr2O72-
D. CrO42- → Cr2O72-
Quiz #5 Application of Cells
1. The corrosion of iron can be prevented by attaching a piece of zinc to the iron because the
A. iron acts as an anode
B. zinc reduces more readily than iron
C. electrons flow from the zinc to the iron
D. 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 anode
B. loses electrons to the cations
C. gains electrons at the cathode
D. gains electrons from the anions
4. An example of electro refining is the
A. extraction of aluminum from bauxite
B. purification of lead from an impure anode
C. recovery of zinc from a zinc sulphide solution
D. production of chlorine from a sodium chloride solution
5. Electroplating always involves the
A. oxidation of anions
B. reduction of cations
C. reduction at the anode
D. oxidation at the cathode
6. Hydrogen and oxygen react to provide energy in a
A. dry cell
B. fuel cell
C. alkaline cell
D. lead-acid storage cell
7. En electrolytic process is used to purify impure lead. The electrodes are
| |ANODE |CATHODE |
|A. |carbon |impure lead |
|B. |pure lead |carbon |
|C. |pure lead |impure lead |
|D. |impure lead |pure lead |
8. In the cell below the half-reaction at the cathode is
A. Cu2+ + 2e- → Cu(s)
B. 2SO42- → S2O82- + 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
| |CATHODE (INERT) |
| | |
| | |
|ANODE (INERT) | |
|Br2 |H2 |
|O2 |Pb |
|Pb |Br2 |
|Br2 |Pb |
10. Under which conditions could an electrochemical cell provide 0.93V?
| |Cathode |
| | |
| | |
|Anode | |
|Cu |Mg |
|Mg |Cu |
|Ag |Pb |
|Pb |Ag |
11. The reduction reaction in the above electrochemical cell is
A. Pb2+ + 2e- → Pb
B. Pb → Pb2+ + 2e-
C. Ag+ + e- → Ag
D. Ag → Ag+ + e-
12. An industrial process involving electrolysis is the reduction of
A. water forming oxygen gas
B. water forming hydrogen gas
C. sea water forming chlorine gas
D. sea water forming bromine liquid
13. To plate a nickel coin with copper
A. the nickel coin must be the cathode
B. the cathode must be made of copper
C. the electrons must flow to the anode
D. 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 fuel
B. electrolyze fuel
C. produce hydrogen
D. 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- → Ni
C. Au → Au3+ + 3e-
D. Au3+ + 3e- → Au
17. Consider the following redox reaction
As2O3 + 2NO3- + 2H2O + 2H+ → 2H3AsO4 + N2O3
In this reaction, nitrogen
A. loses electrons and increases in oxidation number
B. gains electrons and increases in oxidation number
C. loses electrons and decreases in oxidation number
D. gains electrons and decreases in oxidation number
18. In an electrochemical cell, the cathode
A. is reduced
B. loses mass
C. is the reducing agent
D. 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 is
A. 2I- → I2 + 2e-
B. Na+ + e- → Na
C. H2O + ½ O2 + 2H+ + 2e-
D. 2H2O +2e- → H2 + 2OH-
Quiz #6 Corrosion & Cathodic Protection Titration
1. Which of the following metals could be used to cathodically protect a sample of lead?
A. iron
B. gold
C. silver
D. copper
2. A piece of iron can be prevented from corroding by
A. making it a cathode
B. placing it in an acidic solution
C. attaching a small piece of lead to it
D. attaching a small piece of gold to it
3. To determine the [Fe2+] in a solution of FeSO4 by a redox titration, a suitable reagent would be an acidified solution of
A. Cr3+
B. Mn2+
C. SO42-
D. Cr2O72-
4. As a metal corrodes,
A. it gains electrons
B. it becomes reduced
C. it acts as a reducing agent
D. its oxidation number decreases
5. Which method will cathodically protect a piece of iron?
A. Paint the iron
B. Cover the iron with grease
C. Attach a piece of lead tot he iron
D. Attach a piece of magnesium to the iron
6. Corrosion of iron can be prevented by attaching a piece of
A. Mn
B. Cu
C. Pb
D. 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 is
A. less than 0.34 V
B. greater than 1.50 V
C. greater than 0.85 V but less than 1.50 V
D. 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. H2O
D. Na2SO4
9. The reaction that occurs when pieces of lead, zinc, copper and silver are placed in a solution of Ni(NO3)2 is
A. Pb + Ni2+ → Pb2+ + Ni
B. Zn + Ni2+ → Zn2+ + Ni
C. Cu + Ni2+ → Cu2+ + Ni
D. 2Ag + Ni2+ → 2Ag+ + Ni
10. In the electrochemical cell above, the electrons flow from
A. copper to lead through the wire
B. lead to copper through the wire
C. copper to lead through the salt bridge
D. lead to copper through the salt bridge
11. In the electrochemical cell above, the initial Eo value is
A. 0.03 V
B. 0.21 V
C. 0.29 V
D. 0.47 V
12. A reaction that occurs during the corrosion of iron is
A. Fe + 3e- → Fe3+
B. Fe → Fe2+ + 2e-
C. Fe2+ + 2e- → Fe
D. Fe3+ + e- → Fe2+
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. 0.0131 L
B. 0.0262 L
C. 0.200 L
D. 0.400 L
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
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 |
| |
| |
| |
| |
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?
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-
16. Balance the following redox reaction occurring in basic solution:
MnO4- + C2O42- → MnO2 + CO2
17. 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- |
|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+ + 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. 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 |Eo |
|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 | | |
35. 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: | |
|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.
|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.
|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.
|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. Fe
B. CO
C. CO2
D. Fe2O3
2. Consider the following reaction: 2HNO2 + 2I- + 2H+ → 2NO + I2 +2H2O
The oxidation number for each nitrogen atom
A. increases by 1
B. increases by 2
C. decreases by 1
D. 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+ + Cu
D. 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 is
A. Pb + SO42- → 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 V
B. -0.40 V
C. +0.40 V
D. +1.12 V
6. Which of the following involves a nonspontaneous redox reaction?
A. fuel cell
B. electroplating
C. redox titration
D. 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 moles
B. 0.060 moles
C. 1.5 moles
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 + 2H2O
D. 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 V
B. 0.28 V
C. 0.54 V
D. 0.80 V
10. What substances are formed at the anode and cathode during electrolysis of molten sodium chloride?
Anode Cathode
A. 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 is
A. 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- → AgI
B. 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 iron
B. provides an electrolyte
C. acts as a reducing agent
D. acts as an oxidizing agent
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- → Ni
C. Cu → Cu2+ + 2e-
D. Cu2+ + 2e- → Cu
16. The half-reaction that occurs at the cathode is
A. Ni → N2+ + 2e-
B. Ni2+ + 2e- → Ni
C. Cu → Cu2+ + 2e-
D. Cu2+ + 2e- → Cu
17. The cell potential or Eo is
A. 0.41 V
B. 0.78 V
C. 0.34 V
D. 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 Ni
B. through the salt bridge from Cu to Ni
C. through the wire from Cu to Ni
D. through the wire from Ni to Cu
20. Which of the following will not react spontaneously with 1.0 M HCl?
A. tin
B. lithium
C. mercury
D. magnesium
21. Which of the following can be produced by electrolysis from a 1.0 M aqueous solution containing its ion?
A. nickel
B. sodium
C. aluminum
D. 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- → Zn
C. 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- → Zn
D. Zn → Zn2+ + 2e-
25. The corrosion of iron can be prevented by attaching a piece of
A. Mn
B. Cu
C. Pb
D. Sn
26. The oxidation number of carbon in CaC2O4 is
A. +2
B. +3
C. +4
D. +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 copper
C. the electrons must flow to the anode
D. the solution must contain nickel ions
Consider the following electrochemical cell for the next five questions.
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 only
B. I and III only
C. II and III only
D. I, II, and III
29. The balanced equation for the overall reaction is
A. Zn + Cu2+ → Cu + Zn2+
B. Cu + Zn2+ → Zn + Cu2+
C. Zn2+ + Cu → Cu2+ + Zn
D. Cu + Zn → Zn + Cu
30. At equilibrium the voltage of the above cell is
A. -1.10 V
B. 0.00 V
C. +0.42 V
D. +1.10 V
31. This redox reaction occurs because
A. Zn is a stronger oxidizing agent than Cu
B. Zn is a stronger reducing agent than Cu,
C. Cu is a stronger oxidizing agent than Zn
D. Zn2+ is a weaker reducing agent than Cu2+
32. The initial cell voltage at 25 oC is
A. -1.10 V
B. +1.10 V
C. +0.91 V
D. +0.86 V
33. Consider the following redox reaction: Co2+(aq) + 2Ag(s) ⇋ 2Ag+(aq) + Co(s)
The reaction is
A. 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
34. When MnO4- reacts to form Mn2+, the manganese in MnO4- is
A. reduced as its oxidation number increases
B. reduced as its oxidation number decreases
C. oxidized as its oxidation number increases
D. oxidized as its oxidation number decreases
35. The electrolyte used in the alkaline battery is
A. KCl
B. NaOH
C. H2SO4
D. KOH
36. The electrolyte used in an automobile battery is
A. KCl
B. NaOH
C. H2SO4
D. KOH
37. The anode used in the commercial production of Aluminum is
A. C
B. Pt
C. Al
D. Al2O3
38. The anode and cathode used in the electrorefining of impure lead to pure lead are
Anode Cathode
A. Pure Pb Impure Pb
B. Impure Pb Pure Pb
C. Pb2+ Pb
D. Pb Pb2+
39. The anode in the LeClanche or common dry cell is
A. C
B. Zn
C. Mg
D. KOH
40. Which of the following are electrolytic cells
I Electro winning
II Electroplating
III Charging a car battery
IV Fuel cell
A. I and II only
B. I, II, and III only
C. II and II only
D. I, II, III, and IV
Subjective
1. Balance the following in basic solution.
MnO4- + C2O42- → MnO2 + CO2 (basic)
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:
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 24.55
Final 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.
| |
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- → Pb
D 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 V
B -.95 V
C +0.95
D +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
5. Consider the following electrolytic cell:
In the cell above
A I- migrates to the anode and gains electrons
B I- migrates to the cathode and loses electrons
C Na+ migrates to the anode and loses electrons
D 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 electrodes
II A metal being reduced
III A direct current power supply
A I and II only
B I and III only
C II and III only
D 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 is
A 2H+ + 2e- → H2
B 1/2O2 + 2H+ + 2e- → H2O
C 4OH- → O2 + 2H2O + 4e-
D H2 + 2OH- → 2H2O + 2e-
8. 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 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 an aqueous solution of NiI2 using inert electrodes?
A 0.26 V
B 0.28 V
C 0.54 V
D 0.80 V
10.
Which of the following occurs as the cell operates?
A the Zn electrode is reduced and increases in mass
B the Zn electrode is reduced and decreases in mass
C the Zn electrode is oxidized and increases in mass
D 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+ + Zn
C Mg2+ + Ca
D Zn2+ + Mn
12. The concentration of Fe2+(aq) can be determined by a redox titration using
A KBr
B 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 +3
B +4
C +5
D +6
15. Consider the following reaction: 3As2O3 + 4NO3- + 7H2O → 6H3AsO4 + 4NO
The oxidizing agent is
A H+
B H2O
C 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 increased
B reduced since its oxidation number has decreased
C oxidized since its oxidation number has increased
D oxidized since its oxidation number has decreased
17. Consider the following:
I Water
II Oxygen gas
III Nitrogen
At 25oC, a piece of iron rusts in the presence of
A I only
B III only
C I and II only
D II and III only
18. Which of the following represents a redox reaction?
A H2CO3 → H2O + CO2
B CuS + H2 → H2S + Cu
C AgNO3 + NaCl → AgCl + NaNO3
D 2HCl + Na2SO3 → 2NaCl + H2SO3
19. The following reaction occurs in an electrochemical cell:
3Cu2+ + Cr → 2Cr3+ + 3Cu
The Eo for the cell is
A 0.40 V
B 0.75 V
C 1.08 V
D 2.50 V
20. During the corrosion of magnesium, the anode reaction is
A Mg → Mg2+ + 2e-
B Mg2+ + 2e- → Mg
C 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)23- + Br-
23. The process of applying an electric current through a cell to produce a chemical change is called
A corrosion
B ionization
C hydrolysis
D electrolysis
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 strongest to weakest, are
A 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 Cu
C H2S
D Ag2S
26. Consider the redox reaction: 2BrO3- + 10Cl- + 12H+ → Br2 + 5Cl2 + 6H2O
the oxidation half-reaction ivolved in this reaction is
A 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 + 2H2O
D NaOH + HCl → NaCl + H2O
28. During the electrolysis of 1.0 M Na2SO4, the reaction at the cathode is
A Na+ + 1e- → Na
B 2SO42- → 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 → PtO
B PtO3 → PtO2
C Pt(OH)2 → Pt
D Pt(OH)22+ → 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 nickel
B sodium
C aluminum
D magnesium
32. In the electrolysis of molten ZnCl2 using carbon electrodes, the reaction that occurs at the anode is
A Zn → Zn2+ + 2e-
B Zn2+ + 2e- → Zn
C 2Cl- → Cl2 + 2e-
D Cl2 + 2e- → 2Cl-
33. In order for the electrolytic cell to operate, it must have
A a voltmeter
B a salt bridge
C a power supply
D 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.
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.
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
Oxidation and Reduction Reactions Workbook
Notes- double click on the lesson number and download Power Point Viewer if you do not have it.
Worksheets Quiz
1. Oxidation, Reduction, Agents, & Reactions. WS 1
2. Lab: The Strength of Oxidizing Agents.
3. Oxidation Numbers Spontaneous Reactions WS 2 1
4. Oxidation Numbers, Application to Reactions. WS 3
5. Balancing Redox Half Reactions Acid/Base. WS 4 2
6. Balancing Redox Reactions in Acid/Base. WS 5
7. Standard Potentials Using Chart. WS 6 3
8. Electrochemical Cells. WS 7
9. Electrochemical Cells Lab.
10. Electrolytic Cells. WS 8 4
11. Electrolytic Cells Lab.
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
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 - loss of electrons
2. Reduction - gain of electrons
3. Oxidizing agent - causes oxidation by undergoing reduction
4. Reducing agent - causes reduction by undergoing oxidation
Write half reactions for each of the following atoms or ions. Label each as oxidation or reduction.
5. Al -----------> Al3+ + 3e- oxidation
6. S + 2e- ---------> S2- reduction
7. 2O2- ----------> O2 + 4e- oxidation
8. Ba2+ + 2e- -----------> Ba reduction
9. 2N3- ----------> N2 + 6e- oxidation
10. Br2 + 2e- ---------> 2Br- reduction
11. P + 3e- ----------> P3- reduction
12. Ca -----------> Ca2+ + 2e- oxidation
13 Ga3+ + 3e- -----------> Ga reduction
14. S + 2e- ---------> S2- reduction
15. H2 ---------> 2H+ + 2e- oxidation
16. 2H+ + 2e- ---------> H2 reduction
17. 2F- ----------> F2 + 2e- oxidation
18. P3- ----------> P + 3e- oxidation
Balance each spontaneous redox equation. Identify the entities reduced and oxidized. State the reducing agent and the oxidizing agent.
19. Al & Zn2+
2Al + 3Zn2+ → 2Al3+ + 3Zn
oxidized reduced
reducing agent oxidizing agent
20. F2 & O2-
2F2 + 2O2- → 4F- + O2
reduced oxidized
oxidizing agent reducing agent
21. O2 & Ca
2Ca + O2 → 2Ca2+ + 2O2-
oxidized reduced
reducing agent oxidizing agent
22. Al3+ & Li
Al3+ + 3Li → Al + 3Li+
reduced oxidized
oxidizing agent reducing agent
Label the species that is reduced, that is oxidized, the reducing agent and the oxidizing agent.
23. Fe2+ + Co → Co2+ + Fe
Co → Co2+ + 2e- oxidation Fe2+ + 2e- → Fe reduction
24. 3 Ag+ + Ni → Ni3+ + 3 Ag
Ni → Ni2+ + 2e- oxidation Ag+ + 1e- → Ag reduction
25. Cu2+ + Pb → Pb2+ + Cu
Pb → Pb2+ + 2e- oxidation Cu2+ + 2e- → Cu reduction
26. O2 + 2 Sn → O2- + 2 Sn2+
Sn → Sn2+ + 2e- oxidation O2 + 4e- → 2O2- reduction
27. Co2+ + 2 F- → Co + F2
2F- → F2 + 2e- oxidation Co2+ + 2e- → Co reduction
28. List the species (formulas from above) that lose electrons:
Co 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- + Cl2 ----------> Cl- + I2
Substance oxidized I- Reducing agent I-
Oxidizing agent Cl2 Substance reduced Cl2
31. Co + Fe3+ -----------> Co2+ + Fe2+
Substance oxidized Co Reducing agent Co
Oxidizing agent Fe3+ Substance reduced Fe3+
32. Cr6+ + Fe2+ -----------> Cr3+ + Fe3+
Substance oxidized Fe2+ Reducing agent Fe2+
Oxidizing agent Cr6+ Substance reduced Cr6+
Redox Half Reactions and Reactions WS #2
1. State the Oxidation Number of each of the elements that is underlined.
a) NH3 -3 b) H2SO4 6
c) ZnSO3 4 d) Al(OH)3 3
e) Na 0 f) Cl2 0
g) AgNO3 5 h) ClO4- 7
i) SO2 4 j) K2Cr2O4 3
k) Ca(ClO3)2 5 l) K2Cr2O7 6
m) HPO32- 3 n) HClO 1
o) MnO2 4 p) KClO3 5
q) PbO2 4 r) PbSO4 2
s) K2SO4 6 t) NH4+ -3
u) Na2O2 -1 v) FeO 2
w) Fe2O3 3 x) SiO44- -2
y) NaIO3 5 z) ClO3- 5
aa) NO3- 5 bb) Cr(OH)4 4
cc) CaH2 -1 dd) Pt(H20)5(0H)2+ +3
ee) Fe(H2O)63+ +3 ff) CH3COOH 0
2. What is the oxidation number of carbon in each of the following substances?
a) CO 2 b) C 0
c) CO2 4 d) CO32- 4
e) C2H6 -3 f) CH3OH -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) --------> Cu (s) + Zn2+ (aq)
Substance oxidized Zn Substance reduced Cu2+
Oxidizing agent Cu2+ Reducing agent Zn
b) Cl2 (g) + 2 Na (s) --------> 2 Na+ (aq) + 2 Cl- (aq)
Substance oxidized Na Substance reduced Cl2
Oxidizing agent Cl2 Reducing agent Na
WS # 3 Spontaneous and Non-spontaneous Redox Reactions
Describe each reaction as spontaneous or non-spontaneous.
1. Au+3 + Fe+3 -----> Fe+2 + Au nonspontaneous (two oxidizing agents)
2. Pb + Fe+3 ------> Fe+2 + Pb+2 spontaneous
3. Cl2 + F- ------> F2 + 2Cl- nonspontaneous
4. S2O8-2 + 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
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
O2 Cl- Fe Na+
13. Determine the oxidation number for the element underlined.
PbSO4 6 ClO3- 5
HP032- 3 Na2O2 -1
CaH2 -1 Al2(SO4)3 6
NaIO3 5 C4H12 -3
14. Al3+ + Zn ---------> Al + Zn2+
Substance oxidized Zn Oxidizing agent Al3+
15. Cr2O72- + ClO2- ------------> Cr3+ + ClO4-
Substance reduced Cr2O72- Oxidizing agent Cr2O72-
16. State the Oxidation Number of each of the elements that is underlined.
a) NH3 -3 b) H2SO4 6
c) ZnCO3 4 d) Al(OH)3 3
e) Na 0 f) Cl2 0
17. Balance the redox equation using the half reaction method.
Al + 3Ag+ ----------> Al3+ + 3Ag
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 2
20. ClO3- 5
21. HPO32- 3
22. Na202 -1
23. CaH2 -1
24. NaIO3 5
25. C4H12 -3
26. Al2(SO4)3 6
27. Al3+ + Zn ----------> Al + Zn2+
Substance oxidized Zn Oxidizing agent Al3+
28. Cr2O72- + ClO2- ----------------> Cr3+ + ClO4-
Substance reduced Cr2O72- Oxidizing agent Cr2O72-
29. O3 + H2O + SO2 -----> SO42- + O2 + 2H+
Substance oxidized SO2 Reducing agent SO2
30. 3As2O3 + 4NO3- + 7H2O + 4 H+ --------> 6H3AsO4 + 4NO
Substance reduced NO3- Reducing agent As2O3
WS # 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. 5H2O + S2O32- --------------> 2SO42- + 10H+ + 8e-
oxidation
2. 8H+ + 5e- + MnO4- --------------> Mn2+ + 4H2O
reduction
3. 4H2O + As --------------> AsO43- + 8H+ + 5e-
oxidation
4. 7H2O + 2Cr3+ -----------> Cr2O72- + 14H+ + 6e-
oxidation
5. 2H2O + Pb2+ --------------> PbO2 + 4H+ + 2e-
oxidation
6. 8H+ + SO42- + 6e- --------------> S + 4H2O
reduction
7. 4H+ + NO3- + 3e- -------------> NO + 2H2O
reduction
8. 10H+ + 8e- + NO3- --------------> NH4+ + 3H2O
reduction
9. 12H+ + 10e- + 2BrO3- --------------> Br2 + 6H2O
reduction
Balancing Half Cell Reactions
Balance in basic solution.
10. 3e- + 2H2O + NO3- --------------> NO + 4OH-
11. 4H2O + 5e- + MnO4- --------------> Mn2+ + 8OH-
12. 8OH- + As --------------> AsO43- + 4H2O + 5e-
13. 14OH- + 2Cr3+ --------------> Cr2O72- + 7H2O + 6e-
14. 4OH- + Pb2+ --------------> PbO2 + 2H2O + 2e-
15. 4H2O + 6e- + SO42- --------------> S + 8OH-
16. 10 OH- + S2O32- --------------> 2SO42- + 5H2O + 8e-
17. 7H2O + 8e- + NO3- --------------> NH4+ + 10 OH-
18. 6H2O + 10e- + 2BrO3- --------------> Br2 + 12 OH-
19. Determine if each of the following changes is oxidation, reduction or neither.
SO32- --------> SO42- oxidation
CaO --------> Ca reduction
CrO42- --------> Cr2O72- neither
CrO42- --------> Cr3+ reduction
2I- --------> I2 oxidation
IO3- --------> I2 reduction
MnO4- --------> Mn2+ reduction
ClO2- --------> ClO- reduction
20. Cr2O72- + Fe2+ --------> Cr3+ + Fe3+
Substance oxidized Fe2+ Substance reduced Cr2O72-
Oxidizing 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- + 7H2O + 6e- --------------> 14OH- + 2Cr3+
5. NO + 4OH- ------------------> 2H2O + NO3- + 3e-
6. 2H2O + 2e- + SO42- --------------> SO2 + 4OH-
7. 2MnO2 + H2O + 2e- --------------> Mn2O3 + 2OH-
Balance each redox reaction in acid solution using the half reaction method.
8. 8H+ + 3H2O2 + Cr2O72- -------> 3O2 + 2Cr3+ + 7H2O
9. TeO32 - + 2N2O4 + H2O -------> Te + 4NO3- + 2H+
10. 4H+ + 4ReO4- + 7IO- -------> 7IO3- + 4Re + 2H2O
11. 8H+ + 5PbO2 + I2 -------> 5Pb2+ + 2IO3- + 4H2O
12. 12H2O + 8As -------> 3H2AsO4- + 5AsH3 + 3H+
Balance each redox reaction in basic solution using the half reaction method.
13. 3O2 + 8OH- + 2Cr3+ -------> H2O + 3H2O2 + Cr2O72-
14. H2O + Te + 4NO3- -------> TeO32- + 2OH- + 2N2O4
15. 7IO3- + 4OH- + 4Re -------> 4ReO4- + 7IO- + 2H2O
16. 8OH- + 5Pb2+ + 2IO3- -------> 5PbO2 + I2 + 4H2O
17. 7H2O + Cr2O72- + 3Hg -------> 3Hg2+ + 14OH- + 2Cr3+
State of the change represents oxidation, reduction or neither (use oxidation #s).
18. MnO2 --------> Mn2O3 reduction
19. NH3 --------> NO2 oxidation
20. HClO4 -------> HCl + H2O reduction
21. O2 --------> O2- reduction
22. P2O5 --------> P4H10 reduction
Determine the oxidation number
23. H2SO4 6 22. HSO4- 6
24. P4 0 23. NaH -1
25. UO3 6 24. Na2O2 -1
26. U2O5 5 25. PbSO4 2
WS #6 Review
1. Describe each in your own words
1. Oxidation - loss of electrons
2. Reduction - gain of electrons
3. Oxidizing agent - causes oxidation by undergoing reduction
4. Reducing agent - causes reduction by undergoing oxidation
2. Write half reactions for each. Describe as oxidation or reduction. Circle all oxidizing agents.
a) Na -----------> Na+ + e- oxidation
b) Ca -----------> Ca2+ + 2e- oxidation
c) Al3+ + 3e- -----------> Al reduction
d) 2F1- ----------> F2 + 2e- oxidation
e) N2 + 6e- ----------> 2N3- reduction
f) 2O2- ----------> O2 + 4e- oxidation
3. Write the reaction between the following: Use the half reaction method.
a) Ca + Al(NO3)3
3Ca + 2Al3+ -------------> 2Al + 3Ca2+
b) Sn + AgNO3
Sn + 2Ag+ -------------> 2Ag + Sn2+
c) Sn + Au(NO3)3
3Sn + 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+ + 2e- -----------> Ni
Mn2+ + 2e- -----------> Mn
Al3+ + 3e- -----------> 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+ + 1e- -----------> Ag
Pb2+ + 2e- -----------> Pb
Ca2+ + 2e- -----------> 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 + 2e- --------> 2Cl-
Ag+ + 1e- -----------> Ag
Mg2+ + 2e- -----------> 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+ + 2e- -----------> Ni
Mn2+ + 2e- -----------> Mn
Al3+ + 3e- -----------> 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 + 2e- --------> 2Cl-
Br2 + 2e- --------> 2Br-
I2 + 2e- --------> 2I- strongest reducing agent
Classify as oxidation, reduction or neither.
11. SO42- --------> S2- reduction
12. MnO2 --------> MnO4- oxidation
13. Cr2O72- --------> CrO42- neither
14. IO3- --------> I2 reduction
15. Given the following lab data
SnCl2 & Ni Spontaneous
Ni(NO3)2 & Fe Spontaneous
Cr(NO3)3 & Fe Non spontaneous.
i) Write three balanced equations.
Ni + Sn2+ -------------> Ni2+ + Sn
Fe + Ni2+ -------------> Fe2+ + Ni
Fe + Cr3+ Sn
Ni2+ + 2e- -----------> Ni
Fe2+ + 2e- -----------> Fe
Cr3+ + 3e- -----------> 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? No, because Fe2+ is a weaker oxidizing agent than Sn2+
16. 2H+ + 2MnO4- + 5H2S --------> 5S + 6H2O + 2MnO
oxidizing agent reducing agent
17. 2H+ + 10SO42- + 4Br2 ----------> 5S2O32- + 8BrO3- + H2O
oxidizing agent reducing agent
18. Balance in basic solution
2MnO4- + 5H2S --------> 5S + 2MnO + 4H2O + 2OH-
19. Describe as spontaneous or non-spontaneous. Use your reduction potential chart.
a) ZnCl2 & Cu nonspontaneous
b) CuCl2 & NaCl nonspontaneous
c) Br2 & Fe2+ spontaneous
d) H2S & Al3+ nonspontaneous
20. Can you keep HCl in a Zn container? No, Spontaneous reaction.
What about an Au container? Yes, nonspontaneous reaction.
Balance in basic solution
21. H2O + 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. Br2 oxidizing agent 1.09 v
23. Fe2+ oxidizing agent / reducing agent -0.45 v / 0.77 v
24. MnO4- (water) oxidizing agent 0.60 v
25. Ni reducing agent -0.26 v
26. Cr3+ oxidizing agent -0.74 v
27. H2O oxidizing agent / reducing agent -0.40 v / +0.80 v
Indicate as spontaneous or non-spontaneous.
28. MnO4- & Fe2+ non-spontaneous
29. Cu2+ & Br- non-spontaneous
30. HNO3 & Ag spontaneous
31. MnO4- (acid) & H2O spontaneous
32. Ni(s) & Al3+ non-spontaneous
33. HCl & Mg spontaneous
Write each oxidation and reduction half reaction for each question above. Determine the Eo for each. Calculate the Eo for the overall reaction.
34. MnO4- + 2H2O + 3e- --------> MnO2 + 4OH- +0.60 v
3(Fe2+ -----------> Fe3+ + 1e-) -0.77 v
MnO4- + 2H2O + 3Fe2+ -----------> 3Fe3+ + MnO2 + 4OH- -0.17 v
35.
36. NO3- + 4H+ +3e- -----------> NO + 2H2O +0.96 v
3(Ag ----------> Ag+ + 1e-) -0.80 v
NO3- + 4H+ + 3Ag ----------> NO + 2H2O + 3Ag+ +0.16 v
37.
38.
39. 2H+ + 2e- ------> H2 0.00 v
Mg ----------> Mg2+ + 2e- 2.37 v
Mg + 2H+ ----------> Mg2+ + H2 2.37 v
WS # 7 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
12. Ag / AgNO3 ║ H2 / HCl
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
Cathode: Na+ + 1e- → Na(s) -2.71 v Anode: 2Cl- → Cl2 + 2e- -1.36 v
Overall: 2Na+ + 2Cl- → Cl2 + 2Na(s) -4.07 v MTV = +4.07 v
12. Aqueous Na2SO4
Cathode: 2H2O + 2e- → H2 + 2OH- -0.41 v Anode: H2O → 2H+ + 1/2O2 + 2e- -0.82 v
Overall: H2O → H2 + 1/2O2 -1.23 v MTV = +1.23 v
13. Liquid K2O
Cathode: K+ + 1e- → K(s) -2.93 v Anode: 2O2- → O2 + 4e- ? v
Overall: 4K+ + 2O2- → O2 + 4K(s) -? v MTV = +? v
14. 1.0 M LiI
Cathode: Cathode: 2H2O + 2e- → H2 + 2OH- -0.41 v Anode: 2I- → I2 + 2e- -0.54 v
Overall: 2H2O + 2I- → I2 + H2 + 2OH- -0.95 v 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 + 2OH-
0.250L MnO4- x 0.200 mol x 2 mol MnO2 x 86.9g = 4.34g
L 2 mol MnO4- mol
16. Determine the oxidation number for each underlined atom.
MnO2 4 Cr2O7-2 6 IO3- 5 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+ + 3Fe E0 = 1.21 v Spontaneous. There would be a reaction!
19. Draw an electrochemical cell using Cu and Ag electrodes.
Cathode (+) Anode (-)
Ag Cu
Ag+ + 1e---------> Ag 0.80v Cu -------> Cu2 + 2e -0.34v
2Ag+ + Cu ------> 2Ag + Cu2+ E0 = 0.46 v spontaneous
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 2MnO4- = 1.88M
0.100L
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
Oxidizing agent H+ Reducing agent Mg
WS #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: Pb Cathode: Ag
Anode reaction: Pb --------> Pb2+ + 2e- Cathode reaction: Ag+ + 1e- -------> Ag
Overall reaction: Pb + 2Ag+ -----> Pb2+ + 2Ag Voltage: 0.93v
2. ZnCl2(l) electrolytic cell (electro-winning)
Anode: C Cathode: C
Anode reaction: 2Cl- --------> Cl2 + 2e- Cathode reaction: Zn2+ + 2e- -------> Zn
Overall reaction: 2Cl- + Zn2+ -----> Cl2 + Zn MTV: +2.12 v
3. CuSO4(aq) electrolytic cell (electro-winning)
Anode: C Cathode: C
Anode reaction: H2O --------> 2H+ + 1/2O2 + 2e- Cathode reaction: Cu2+ + 2e- -------> Cu
Overall reaction: H2O + Cu2+ -----> 2H+ + 1/2O2 + Cu MTV: +0.48 v
4. The electrolysis of 1M NaI (electro-winning)
Anode: C Cathode: C
Anode reaction: 2I- --------> I2 + 2e- Cathode reaction: 2H2O + 2e- -------> H2 + 2OH-
Overall reaction: 2H2O + 2I- -----> 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: C Cathode: C
Anode reaction: 2O2- -------> O2 + 4e- Cathode reaction: Al3+ + 3e- -------> Al
Overall reaction: 6O2- + 4Al3+ -----> 3O2 + 4Al
6. The reaction needed to electroplate a copper penny with silver.
Anode: Ag Cathode: penny
Anode reaction: Ag-----> Ag+ + e- Cathode reaction: Ag+ + e- -----> Ag
7. The reaction needed to nickel plate a copper penny.
Anode: Ni Cathode: penny
Anode reaction: Ni-----> Ni+2 + 2e- Cathode reaction: Ni2+ + 2e- -----> Ni
Possible Electrolyte Ni(NO3)2
8. The reaction used in the electrorefining of lead.
Anode: Impure Lead Cathode: Pure Lead
Anode reaction: Pb-----> Pb+2 + 2e- Cathode reaction: 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 Cathode: Zn
Anode reaction: Mg --------> Mg2+ + 2e- Cathode reaction: Zn+2 + 2e- -------> Zn
Overall reaction: Mg + Zn2+ -----> Mg2+ + Zn Voltage: 1.61v
2. The electrolytic cell used to produce Al.
Electrolyte: Al2O3 Phase (aqueous or molten) Molten
Anode: C Cathode: C
Anode reaction: 2O2- -------> O2 + 4e- Cathode reaction: Al3+ + 3e- -------> Al
Overall reaction: 6O2- + 4Al3+ -----> 3O2 + 4Al
3. The electrolysis KI(aq)
Anode: C Cathode: C
Anode reaction: 2I- --------> I2 + 2e- Cathode reaction: 2H2O + 2e- -------> H2 + 2OH-
Overall reaction: 2H2O + 2I- -----> H2 + 2OH- + I2 MTV: +0.95 v
4. The electrorefining of Pb
Anode: Impure Lead Cathode: Pure Lead
Anode reaction: Pb-----> Pb+2 + 2e- Cathode reaction: Pb2+ + 2e- -----> Pb
5. Nickel plating an iron nail.
Anode: Ni Cathode: nail
Anode reaction: Ni-----> Ni+2 + 2e- Cathode reaction: Ni2+ + 2e- -----> Ni
Possible Electrolyte Ni(NO3)2 The -ve side of the power supply is connected to the nail
6. Draw an Ag/ Zn electrochemical cell.
Anode: Zn Cathode: Ag
Anode reaction: Zn --------> Zn2+ + 2e- Cathode reaction: Ag+ + 1e- -------> Ag
Overall reaction: Zn + 2Ag+ -----> Zn2+ + 2Ag Voltage: 1.56v
7. Draw a KF(l) electrolytic cell.
Anode: C Cathode: C
Anode reaction: 2F- --------> F2 + 2e- Cathode reaction: K+ + e- -------> K
Overall reaction: 2F- + 2K+ -----> Cl2 + K MTV: +5.80v
8. Draw a KF(aq) electrolytic cell.
Anode: C Cathode: C
Anode reaction: H2O --------> 2H+ + 1/2O2 + 2e- Cathode reaction: 2H2O + 2e- -------> H2 + 2OH-
Overall reaction: H2O -----> H2 + 1/2O2 MTV: +1.23 v
9. Draw a FeI2(aq) electrolytic cell.
Anode: C Cathode: C
Anode reaction: 2I- --------> I2 + 2e- Cathode reaction: Fe2+ + 2e- -------> Fe
Overall reaction: Fe2+ + 2I- -----> 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 Cathode: Cd
Anode reaction: Pb --------> Pb2+ + 2e- 0.13v Cathode reaction: Cd+2 + 2e- -------> Zn x volts
Overall reaction: Pb + Cd2+ -----> Pb2+ + Cd Voltage: 0.50v
0.13 + x = 0.50 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 |electrolyte |
|Cl2 production |C |2Cl- ------> Cl2 |C |Na+ + e- -----> Na |NaCl(l) |
| | |+ 2e- | | | |
|Leclanche or Common |Zn |Zn-->Zn+2 + 2e- |C/MnO2 |Mn+4 +1e- -----> Mn+3 |NH4Cl and MnO2 |
|Dry Cell | | | | | |
|Nickel Plating |Ni |Ni-->Ni+2 + 2e- |Metal to be |Ni2+ +2e- -----> Ni |Ni(NO3)2 |
| | | |plated | | |
|Lead Storage or Car |Pb |Pb ---> Pb+2+ 2e- |PbO2 |PbO2 + SO4-2 + 4OH-1 + 2e- -----> |H2SO4 |
|Battery | | | |PbSO4 + 2H2O | |
|Fuel Cell |C |H2 + 2OH- ---> 2H2O +|C |O2 + 2H2O +4e-----> 4OH- |KOH |
| | |2e- | | | |
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
-----------------------
| 1.0 M KNO3 |
|1 M Mg(NO3)2 |
|Ag |
|Mg |
|1 M AgNO3 |
| |
|Pt |
| |
|Pt |
|1 M CuSO4 |
| 1.0 M KNO3 |
|1M Zn(NO3)2 |
| |
| |
|Pb |
|Zn |
|1M Pb(NO3)2 |
| voltmeter |
|Power Source |
|- + |
| |
|Pt |
| |
|Pt |
|Molten MgCl2 |
| 1.0 M KNO3 |
|1 M Mg(NO3)2 |
| |
| |
|Cu |
|Mg |
|1 M CuSO4 |
| voltmeter |
| 1.0 M KNO3 |
|1 M Ni(NO3)2 |
| |
| |
|Ag |
|Ni |
|1 M AgNO3 |
| voltmeter |
| 1.0 M KNO3 |
|1 M Au(NO3)3 |
| |
| |
|Pb |
|Au |
|1 M Pb(NO3)2 |
| voltmeter |
|Power Source |
|- + |
| |
|Pt |
| |
|Iron Key |
|1.0 M CuSO4 |
A.
B.
C.
D.
A.
B.
C.
D.
| 1.0 M KNO3 |
|1M Pb(NO3)2 |
| |
| |
|Cu |
|Pb |
|1M Cu(NO3)2 |
| voltmeter |
| 1.0 M KNO3 |
|1 M Zn(NO3)2 |
| |
| |
|Cu |
|Zn |
|1 M Cu(NO3)2 |
| voltmeter |
|1 M HCl |
| |
|Cu |
|1 M Cu(NO3)2 |
| voltmeter |
|H2(g) |
| 1.0 M KNO3 |
| |
|Power Source |
|- + |
| |
|C |
| |
|C |
|Molten Al2O3 |
|Power Source |
|- + |
| |
|Cu |
| |
|Cu |
|1 M NaF |
| |
| |
|Ni |
|Cuu |
| 1.0 M KCl |
|1M Cu(NO3)2 |
|1M Ni(NO3)2 |
| voltmeter |
| |
| |
|Zn |
|Cuuu |
| 1.0 M KNO3 |
|1M Cu(NO3)2 |
|1M Zn(NO3)2 |
| voltmeter |
|Power Source |
| |
|Impure Pb |
| |
|Pure Pb |
voltmeter
| |
| |
|Mn |
|Snnnn |
| 1.0 M KNO3 |
|1M Sn(NO3)2 |
|1M MnNO3)2 |
| |
| |
|Ni |
|Pd |
| 1.0 M KCl |
|1M Pd(NO3)2 |
|1M Ni(NO3)2 |
| voltmeter |
|Power Source |
|+ - |
|Inert |
|Electrode |
|Inert |
|Electrode |
|Molten NaI(l) |
| |
| |
|Ni |
|Zn |
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|1.0 M KNO3 |
|1M Zn(NO3)2 |
|1M Ni(NO3)2 |
| voltmeter |
|Power Source |
| |
|Impure Pb |
| |
|Pure Pb |
| |
| 1.0 M KNO3 |
|1 M Zn(NO3)2 |
| |
| |
|Cu |
|Zn |
|1 M Cu(NO3)2 |
| voltmeter |
|Zn → Zn2+ + 2e- |
|oxidation |
|anode |
|0.76 v |
|loses mass |
|Cu has greater electron affinity |
|Cu2+ + 2e- → Cu |
|reduction |
|cathode |
|0.34 v |
|gains mass |
| 2 e- |
| 2 e- |
| |
| |
|Zn2+ |
| |
|NO3- |
| |
|Cu2+ |
| |
|NO3- |
| NO3- K+ |
| |
|Cu2+ + Zn → Zn2+ + Cu 1.10 v |
| 1.0 M KNO3 |
|1 M HCl |
| |
| |
|Cu |
|H2 |
| |
|1 M Ag(NO3)2 |
| voltmeter |
|H2 → 2H+ + 2e- |
|oxidation |
|anode |
|0.00 v |
| |
|Ag has a greater electron affinity |
|2Ag+ + 2e- → 2Ag |
|reduction |
|cathode |
|0.80 v |
|gains mass |
| 2 e- |
| 2 e- |
| |
| |
|H+ |
| |
|Cl- |
| |
|Ag+ |
| |
|NO3- |
| NO3- K+ |
| |
|2Ag+ + H2 → 2Ag + 2H+ 0.80 v |
|Power Source |
|- + |
| |
|Pt |
| |
|Pt |
|Na+ |
|Cl- |
|Power Source |
|- + |
| |
|C |
| |
|C |
|Na+ |
|SO42-|
|H2O |
|Power Source |
|- + |
| |
|Pt |
| |
|Pt |
|K+ |
|O2- |
|Power Source |
|- + |
| |
|Pt |
| |
|Pt |
|Li+ |
|I- |
................
................
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