Oxidation and Reduction Reactions Workbook 5/Workbook... · Web viewElectrochemical Cells Lab. 10. Electrolytic Cells. WS 8 4 11. Electrolytic Cells Lab. 12. Application of Electrochemical

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. 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-

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. 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+

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


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 __________

28. Cr2O7

2- + 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

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.

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

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

- → TeO32- + N2O4

15. IO3

- + Re → ReO4- + IO-

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

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. 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.

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

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. 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 + 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. 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)


4. The electrolysis of 1M NaI (electrowinning)


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.


8. The reaction used in the electrorefining of lead.


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 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?

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.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 oxidizedB. is called the cathode

1.0 M KNO3

1 M Mg(NO3)2

AgMg

1 M AgNO3

Pt

Pt

1 M CuSO4

C. gains mass at the electrodeD. 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, AuB. Pt, Au, AgC. Au, Ag, PtD. 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 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 of

A. 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 is

A. FeB. Cu2+

C. H2OD. SO4

2-

9. A substance is oxidized when it

A. 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 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 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 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 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 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+ + 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-reaction HClO2 ⇄ HClO The balanced half-reaction would have

A. 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)4

2+ is

A. +2B. 0C. +4D. +1/2

9. Consider the following half-reaction: BrO- → Br- (basic)

The balanced equation for the half-reaction is

A. 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+ + 3H2O

The species that loses the electron is

A. H2OB. MnO4

-

C. CH3CHOD. CH3COOH

11. Hydrogen has an oxidation number of –1 in

A. H2

B. NaHC. H2OD. KOH

12. Consider the following:

2NO3- + 4H+ + 2e- → N2O4 + 2H2O

This equation represents

A. reductionB. oxidationC. neutralizationD. 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. AlB. MnO4

-

C. Al(OH)4-

D. MnO2


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 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 undergoes

A. reductionB. oxidationC. electrolysisD. neutralization

19. The oxidation number in carbon in CaC2O4 is

A. +2B. +3C. +4D. +6


2Cr3+(aq) + 3Cl2(aq) + 7H2O(l) → Cr2O7

2-(aq) + 6Cl-

(aq) + 14H+(aq)

The species which loses electrons is

A. Cl2

B. Cr3+

C. H2OD. Cr2O7

2-

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 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 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 METAL1 CuCl2 Zn2 CuCl2 Ag3 CuCl2 Ni

Reactions occur in

A. beaker 2 onlyB. beakers 1, 2, and 3C. beakers 1 and 2 onlyD. beakers 1 and 3 only


3SO2 + 3H2O + ClO3- → 3SO4

2- + 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 VB. +0.39 VC. +1.21 VD. +1.54 V

9. A solution of 1.0 M Co(NO3)2 should be stored in a container made of

A. 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 are

A. IO3- and I2

B. SO42- and S

C. BrO3- and Br -

D. AuCl4- and Au


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. ZnB. 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 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 is

A. -0.89 VB. -0.63 VC. +0.63 VD. +0.89 V

3. In an operating lead-zinc electrochemical cell shown above, the cathode

A. gains mass as anions are reducedB. loses mass as anions are reducedC. gains mass as cations are reducedD. loses mass as cations are reduced

1.0 M KNO3

1M Zn(NO3)2

PbZn

1M Pb(NO3)2

voltmeter

4. The equation for the half-reaction at the anode is

A. 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 from

A. 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 8. As this cell operates

A. 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 migrate

A. towards the anode through the wireB. towards the cathode through the wire

Power Source

- +

Pt

Pt

Molten MgCl2

C. towards the anode through the salt bridgeD. towards the cathode through the salt bridge

10. As the above electrochemical cell operates

A. 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 is

A. 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 electrolysis

A. 1.0 M KIB. 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- → ZnD. Zn → Zn2+ + 2e-

1.0 M KNO3

1 M Mg(NO3)2

CuMg

1 M CuSO4

voltmeter

15. The initial cell voltage at 25oC is

A. -1.06 VB. -0.54 VC. +0.54 VD. +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)

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

18. The direction of the electron flow is

A. 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 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 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-

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 anodeB. zinc reduces more readily than iron

C. 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

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

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.B.C.D.

A. Ni → Ni2+ + 2e-

B. Ni2+ + 2e- → NiC. 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 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 cathode

A. 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 is

A. 2I- → I2 + 2e-

B. Na+ + e- → NaC. 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. ironB. goldC. silverD. copper

2. A piece of iron can be prevented from corroding by

A. making it a cathodeB. placing it in an acidic solutionC. 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 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 ironB. 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 of

A. 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 is

A. 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 is

A. 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 from

A. 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

1.0 M KNO3

1M Pb(NO3)2

CuPb

1M Cu(NO3)2

voltmeter

11. In the electrochemical cell above, the initial Eo value is

A. 0.03 VB. 0.21 VC. 0.29 VD. 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- → FeD. 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 LB. 0.0262 LC. 0.200 LD. 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- + 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-

20. How is the breathalyzer reaction used to determine BAC? Write the reaction and describe how it works.

21. 2H+ + Mg → Mg2+ + H2

Determine the Oxidizing agent__________ and the Reducing agent_________

22. Choose a suitable redox reactant to oxidize Cl- toClO4- in a redox titration.

23. Describe as an electrochemical or electrolytic cell:

a) Fuel cell

b) Charging a car battery

c) Discharging a car battery

d) Ni plating

e) Industrial Al production

f) Cl2 production

g) Electrowinning

24) Which of the reactants is gaining electrons? Which of the reactants is the oxidizing agent?

Br2 + SO2 + Na2SO4 + H2O → 2H2SO4 + 2NaBr

25) A student studied the following reactions and she recorded:

Pd2+ + Cu → Pd + Cu2+ spontaneous

Pd2+ + Au → no reaction

Pd2+ + Hg → no reaction

Au3+ + Hg → Au + Hg2+ spontaneous

List the oxidizing agents from strongest to weakest. List the reducing agents from strongest to weakest. Predict if the reaction will occur.

Au3+ + 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.


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

1.0 M KNO3

1 M Zn(NO3)2

CuZn

1 M Cu(NO3)2

voltmeter

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.

1 M HCl

Cu

1 M Cu(NO3)2

voltmeter

H2(g)

1.0 M KNO3

Anode Reaction

Cathode Reaction

Chemicals produced at the anode

Chemicals produced at the cathode

The electrons flow from __to __

The chemical used to lower the mp is:

Which electrode is the anode ?

38. Completely analyze the following electrolytic cell. Note that the electrodes are not inert and because of that, the anode might oxidize.

Power Source

- +

C

C

Molten Al2O3

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. FeB. COC. CO2

D. Fe2O3

Power Source

- +

Cu

Cu

1 M NaF

2. Consider the following reaction: 2HNO2 + 2I- + 2H+ → 2NO + I2 +2H2O

The oxidation number for each nitrogen atom

A. 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 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 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


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 moles

C. 1.5 molesD. 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 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 ironB. provides an electrolyteC. 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- → 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 Ni

NiCuu

1.0 M KCl

1M Cu(NO3)2 1M Ni(NO3)2

voltmeter

C. 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

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 of

A. 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. 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 is

A. 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 V

ZnCuuu

1.0 M KNO3

1M Cu(NO3)2 1M Zn(NO3)2

voltmeter

D. +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 is

A. 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 36. The electrolyte used in an automobile battery is

A. 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 Cathode

A. 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) 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.

Power Source

Impure PbPure Pb

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 24.55Final 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

Chemistry 12 Electrochemistry Practice Test 2

NiPd

1.0 M KCl

1M Pd(NO3)2 1M Ni(NO3)2

voltmeter

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-

cellD 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 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 is

A 2H+ + 2e- → H2

B 1/2O2 + 2H+ + 2e- → H2O

Power Source+ -

Inert Electrode

Inert Electrode

Molten NaI(l)

C 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 an aqueous solution of NiI2

using inert electrodes?

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 using

A 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)

NiZn

1.0 M KNO3

1M Zn(NO3)2 1M Ni(NO3)2

voltmeter

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 is

A H+

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 Nitrogen

At 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 is

A 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 hydrolysisD 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 CuC H2SD Ag2S

26. Consider the redox reaction: 2BrO3

- + 10Cl- + 12H+ → Br2 + 5Cl2 + 6H2Othe 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 + 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

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- → 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.

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:

Power Source

Impure Pb

Pure Pb

a) On the diagram above, label the anode and cathode. b) Write the formula for a suitable electrolyte. c) Write the equation for the reduction half-reaction. d) Write the anode reaction

Oxidation and Reduction Reactions Workbook

Notes- double click on the lesson number and download Power Point Viewer if you do not have it.

Worksheets 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!!

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http://www.mouatonline.com/Teachers/MIannone/Science/Chemistry/Iannone/Redox/Powerpoint/1.IntroductionRedox.ppt

Redox Half Reactions and Reactions WS #1

Define each 1. Oxidation - loss of electrons2. Reduction - gain of electrons3. Oxidizing agent - causes oxidation by undergoing reduction4. 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+ + 3Znoxidized reducedreducing agent oxidizing agent

20. F2 & O2-

2F2 + 2O2- → 4F- + O2

reduced oxidizedoxidizing agent reducing agent

21. O2 & Ca

2Ca + O2 → 2Ca2+ + 2O2-

oxidized reducedreducing agent oxidizing agent

22. Al3+ & Li

Al3+ + 3Li → Al + 3Li+ reduced oxidizedoxidizing 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-

→ Coreduction

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 CoOxidizing 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


a) NH3 -3 b) H2SO4 6c) ZnSO3 4 d) Al(OH)3 3e) Na 0 f) Cl2 0g) AgNO3 5 h) ClO4

- 7i) SO2 4 j) K2Cr2O4 3k) Ca(ClO3)2 5 l) K2Cr2O7 6m) HPO3

2- 3 n) HClO 1o) MnO2 4 p) KClO3 5q) PbO2 4 r) PbSO4 2s) K2SO4 6 t) NH4

+ -3u) Na2O2 -1 v) FeO 2w) Fe2O3 3 x) SiO4

4- -2 y) NaIO3 5 z) ClO3

- 5aa) NO3

- 5 bb) Cr(OH)4 4cc) 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 0c) CO2 4 d) CO3

2- 4e) 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 agentZn

b) Cl2 (g) + 2 Na (s) --------> 2 Na+ (aq) + 2 Cl- (aq)

Substance oxidized Na Substance reduced Cl2

Oxidizing agent Cl2 Reducing agentNa

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 6NaIO3 5 C4H12 -3

14. Al3+ + Zn ---------> Al + Zn2+

Substance oxidized Zn Oxidizing agent Al3+

15. Cr2O72- + ClO2

- ------------> Cr3+ + ClO4-

Substance reduced Cr2O7

2- Oxidizing agent Cr2O72-


a) NH3 -3 b) H2SO4 6c) ZnCO3 4 d) Al(OH)3 3e) Na 0 f) Cl2 0

17. Balance the redox equation using the half reaction method.

Al + 3Ag+ ----------> Al3+ + 3Ag


O2 Cl- Fe Na+

Determine the oxidation number for the element underlined.19. PbSO4 220. ClO3

- 521. HPO3

2- 3

22. Na202 -123. CaH2 -124. NaIO3 525. C4H12 -326. Al2(SO4)3 6 27. Al3+ + Zn ----------> Al + Zn2+

Substance oxidized Zn Oxidizing agent Al3+

28. Cr2O7

2- + ClO2- ----------------> Cr3+ + ClO4

-

Substance reduced Cr2O7

2- Oxidizing agentCr2O7

2-

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 + S2O3

2- --------------> 2SO42- + 10H+ + 8e-

oxidation 2. 8H+ + 5e- + MnO4

- --------------> Mn2+ + 4H2O

reduction 3. 4H2O + As --------------> AsO4

3- + 8H+ + 5e-

oxidation 4. 7H2O + 2Cr3+ -----------> Cr2O7

2- + 14H+ + 6e-

oxidation

5. 2H2O + Pb2+ --------------> PbO2 + 4H+ + 2e-

oxidation 6. 8H+ + SO4

2- + 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 --------------> AsO4

3- + 4H2O + 5e-

13. 14OH- + 2Cr3+ --------------> Cr2O7

2- + 7H2O + 6e-

14. 4OH- + Pb2+ --------------> PbO2 + 2H2O + 2e-

15. 4H2O + 6e- + SO4

2- --------------> S + 8OH-

16. 10 OH- + S2O3

2- --------------> 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- --------> SO4

2- oxidation

CaO --------> Ca reduction

CrO42- --------> Cr2O7

2- neither

CrO42- --------> Cr3+ reduction

2I- --------> I2 oxidation

IO3- --------> I2 reduction

MnO4- --------> Mn2+ reduction

ClO2- --------> ClO- reduction

20. Cr2O7

2- + Fe2+ --------> Cr3+ + Fe3+

Substance oxidized Fe2+ Substance reducedCr2O7

2-

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. Cr2O7

2- + 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 + Cr2O7

2-

14. H2O + Te + 4NO3- ------->TeO3

2- + 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 number23. H2S O4 6 22. HS O4

- 624. P 4 0 23. NaH -125. U O3 6 24. Na2O 2 -1 26. U 2O5 5 25. P b SO4 2

WS #6 Review 1. Describe each in your own words 1. Oxidation - loss of electrons2. Reduction - gain of electrons3. Oxidizing agent - causes oxidation by undergoing reduction4. 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- oxidationb) Ca -----------> Ca2+ + 2e- oxidationc) Al3+ + 3e- -----------> Al reductiond) 2F1- ----------> F2 + 2e- oxidatione) N2 + 6e- ----------> 2N3- reductionf) 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- -----------> MnAl3+ + 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- -----------> PbCa2+ + 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- -----------> AgMg2+ + 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- -----------> MnAl3+ + 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. SO4

2- --------> S2- reduction12. MnO2 --------> MnO4

- oxidation13. Cr2O7

2- --------> CrO42- neither

14. IO3- --------> I2 reduction

15. Given the following lab data

SnCl2 & Ni SpontaneousNi(NO3)2 & Fe SpontaneousCr(NO3)3 & Fe Non spontaneous.

i) Write three balanced equations.Ni + Sn2+ -------------> Ni2+ + SnFe + Ni2+ -------------> Fe2+ + NiFe + Cr3+ <------------- Fe2+ + Cr

ii) Rank the oxidizing agents in decreasing order of strength.

strongest oxidizing agent Sn2+ + 2e- -----------> SnNi2+ + 2e- -----------> NiFe2+ + 2e- -----------> FeCr3+ + 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+ + 10SO4

2- + 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 nonspontaneousb) CuCl2 & NaCl nonspontaneousc) Br2 & Fe2+ spontaneousd) 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 + 10SO4

2- + 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 v22. Br2 oxidizing agent 1.09 v23. Fe2+ oxidizing agent / reducing agent -0.45 v / 0.77 v24. MnO4

- (water) oxidizing agent 0.60 v25. Ni reducing agent -0.26 v26. Cr3+ oxidizing agent -0.74 v27. 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-spontaneous30. HNO3 & Ag spontaneous31. MnO4

- (acid) & H2O spontaneous32. Ni(s) & Al3+ non-spontaneous33. 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 v3(Fe2+ -----------> Fe3+ + 1e-) -0.77 v

MnO4- + 2H2O + 3Fe2+ -----------> 3Fe3+ + MnO2 + 4OH- -0.17 v

35. 36. NO3

- + 4H+ +3e- -----------> NO + 2H2O +0.96 v3(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 vMg + 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

1.0 M KNO3

1 M Zn(NO3)2

CuZn

1 M Cu(NO3)2

voltmeter

Zn → Zn2+ + 2e-

oxidationanode0.76 vloses mass

Cu has greater electron affinityCu2+ + 2e- → Cu reductioncathode0.34 vgains 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

CuH2

1 M Ag(NO3)2

voltmeter

H2 → 2H+ + 2e-

oxidationanode0.00 v

Ag has a greater electron affinity2Ag+ + 2e- → 2Ag reductioncathode0.80 vgains mass

2 e- 2 e-

H+

Cl-

Ag+

NO3-

NO3- K+

2Ag+ + H2 → 2Ag + 2H+ 0.80 v

WS # 8 1. In an electrolytic cell, reduction occurs at the negative electrode and oxidation occurs at the positive electrode. 2. If there are two possible reduction reactions, the highest one on the chart occurs. 3. For reduction, the chart is read from left to right. 4. For oxidation, the chart is read from right to left and the sign of the voltage is changed. 5. If there are two possible oxidation reactions, the lowest one on the chart occurs. 6. Corrosion of a metal is oxidation. 7. Electrolysis uses electrical energy. 8. Electrochemical cells produce electrical energy. 9. Electrolytic cells use electrical energy. 10. What is the standard reference cell? hydrogen Eo = O v Draw and completely analyze each electrolytic cell.11. Molten NaCl

Power Source

- +

Pt

Pt

Na+

Cl-

Cathode: Na+ + 1e- → 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

Power Source

- +

C

C

Na+

SO42-

H2O

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

Power Source

- +

Pt

Pt

K+

O2-

Power Source

- +

Pt

Pt

Li+

I-

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 electricityAnode- an electrode that is the site of oxidationCathode- an electrode that is the site of reduction Spontaneous- a reaction that occurs naturally and has a positive voltageElectron 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 CuAg+ + 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 2MnO 4- = 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 Cr2O7

2- and ethanol C2H5OH. If alcohol is present in your breath sample, it will react with a solution of Cr2O7

2- 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.



Zn Zn-->Zn+2 + 2e- C Mn+4 +1e- -----> Mn+3 NH4Cl and MnO2

Alkaline Cell Zn Zn-->Zn+2 + 2e- C Mn+4 +1e- -----> Mn+3 KOH and MnO2

Lead Storage or Car Battery Pb Pb ---> Pb+2+ 2e- PbO2

PbO2 + HSO4- + 3H+ + 2e-

-----> PbSO4 + 2H2O H2SO4

Fuel Cell C H2 + 2OH- ---> 2H2O + 2e- C ½O2 + H2O +2e-----> 2OH- KOH

12. 2HIO3 + 5H2SO3 ----------> I2 + 5H2SO4 + H2O oxidizing agent HIO3 substance oxidized H2SO3 substance reduced HIO3 reducing agent H2SO3 13.What is the electrolyte in a fuel cell? KOH 14. What is the fuel in a fuel cell? H2 and O2

15. Describe the differences and similarities between an electrolytic and electrochemical cell. Electrolytic Electrochemical Uses electricity Produces electricityNonspontaneous SpontaneousMakes chemicals Uses chemicalsInert carbon electrodes Usually has a salt bridgeThe negative electrode is reduction The higher metal is reduction Oxidation occurs at the anode and reduction occurs at the cathode.Anions migrate to the anode and cations migrate to the cathode.Electrons go from anode to cathode through the wire. 16. Describe and give two examples of electrowinning. The electrolysis of water to make H2 and O2. The electrolysis of Al2O3 to make Al and O2. 17. Describe and give one example of electrorefinning. The electrorefinning of Pb. 18. List three metals that can be won from aqueous solution. Pb Au Ag Zn

Cu Fe Sn 19. List three metals that cannot be won from aqueous solution. Na K Li Ca Mg

Al 20. What is the electrolyte in a fuel cell, alkaline battery, Dry Cell (Leclanche) and lead acid battery? KOH KOH & MnO2 NH4Cl & MnO2 PbSO4

21. State two metals that can be used to cathodically protect Fe. Describe how they protect iron from corrosion. Zn and Mg. When attached to Fe they form an electrochemical cell. Zn or Mg is a stronger reducing agent (lower on the chart) and is the anode and Fe is the cathode. Since the cathode is the site of reduction, Fe cannot oxidize or corrode. 22. Write the half reaction that describes the corrosion of iron. Fe --------> Fe2+ +2e-

23. Write the half reaction that describes the reduction reaction that occurs when iron corrodes in air and water. 2e- + H2O + 1/2O2 ----------> 2OH-

24. Why does iron corrode faster in salt water? The salt acts like a salt-bridge and increases the rate of reaction in an electrochemical cell. 25. Write the anode and cathode reaction in an electrolytic cell with a CaCl2 (l) electrolyte. Cathode: Ca2+ + 2e- ---------> Ca Anode: 2Cl- ----------> Cl2 + 2e- 26. Explain why you would choose Zn or Cu to cathodically protect iron? Zn. It is a stronger reducing agent than Fe and it will allow Fe to be the cathode, which cannot corrode. 27. Choose a suitable redox reactant to oxidize Cl- to ClO4

- in a redox titration. MnO4

- in acid gives a spontaneous reaction as well as a color change from purple to clear. 28. Describe as an electrochemical or electrolytic cell:

a) Fuel cell electrochemical b)Charging a car batteryelectrolyticc) Discharging a car battery electrochemical d) Ni platingelectrolytice) Industrial Al production electrolytic f) Cl2 productionelectrolytic

29) Write the anode and cathode reactions.


Cl2 production C 2Cl- ------> Cl2 + 2e- C Na+ + e- -----> Na NaCl(l)


Zn Zn-->Zn+2 + 2e- C/MnO2 Mn+4 +1e- -----> Mn+3 NH4Cl and MnO2

Nickel Plating Ni Ni-->Ni+2 + 2e- Metal to be plated Ni2+ +2e- -----> Ni Ni(NO3)2

Lead Storage or Car Battery Pb Pb ---> Pb+2+ 2e- PbO2

PbO2 + SO4-2 + 4OH-1 + 2e-

-----> PbSO4 + 2H2O H2SO4

Fuel Cell C H2 + 2OH- ---> 2H2O + 2e- C O2 + 2H2O +4e-----> 4OH- KOH

30) Al and AgNO3(aq) are mixed and the surface of the Al darkens. List the two oxidizing agents in decreasing strength. List the two reducing agents in decreasing strength. Oxidizing Agents Ag+ Al3+

Reducing Agents Al Ag

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Oxidation and Reduction Reactions Workbook 5/Workbook... · Web viewElectrochemical Cells Lab. 10. Electrolytic Cells. WS 8 4 11. Electrolytic Cells Lab. 12. Application of Electrochemical