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. N^3-

10. Br₂

11. P

12. Ca

13 Ga³⁺

14. S

15. H₂

16. H⁺

17. F^-

18. P^3-

Balance each spontaneous redox equation. Identify the entities reduced and oxidized. State the reducing agent and the oxidizing agent.

19. Al & Zn²⁺

20. F₂ & O^2-

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 ⇄ O^2- + 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) HPO₃^2- _____ n) HClO _____

o) MnO₂ _____ p) KClO₃ _____

q) PbO₂ _____ r) PbSO₄ _____

s) K₂SO₄ _____ t) NH₄⁺ _____

u) Na₂O₂ _____ v) FeO _____

w) Fe₂O₃ _____ x) SiO₄^4- _____

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) CO₃^2- _____

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) Cl_{2 (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 → 2SO₄^2- + 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^- __________

HP0₃^2- __________ Na₂O₂ __________

CaH₂ __________ Al₂(SO₄)₃ __________

NaIO₃ __________ C₄H₁₂ __________

14. Al³⁺ + Zn → Al + Zn²⁺

Substance oxidized _______ Oxidizing agent ________

15. Cr₂O₇^2- + ClO₂^- → Cr³⁺ + ClO₄^-

Substance reduced ________ Oxidizing agent ________

16. State the Oxidation Number of each of the elements that is underlined.

a) NH₃ __________ b) H₂SO₄ __________

c) ZnCO₃ __________ d) Al(OH)₃ __________

e) Na __________ f) Cl₂ __________

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

Al & AgNO₃

18. Circle each formula that is able to lose an electron

O₂ Cl^- Fe Na⁺

Determine the oxidation number for the element underlined.

19. PbSO₄ __________

20. ClO₃^- __________

21. HPO₃^2- __________

22. Na₂0₂ __________

23. CaH₂ __________

24. NaIO₃ __________

25. C₄H₁₂ __________

26. Al₂(SO₄)₃ __________

27. Al³⁺ + Zn → Al + Zn²⁺

Substance oxidized __________ Oxidizing agent __________

28. Cr₂O₇^2- + ClO₂^- → Cr³⁺ + ClO₄^-

Substance reduced __________ Oxidizing agent __________

29. O₃ + H₂O + SO₂ → SO₄^2- + 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- → SO₄^2-

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- → SO₄^2-

17. NO₃^- → NH₄⁺

18. BrO₃^- → Br₂

19. Determine if each of the following changes is oxidation, reduction or neither.

SO₃^2- → SO₄^2- ________________

CaO → Ca ________________

CrO₄^2- → Cr₂O₇^2- ________________

CrO₄^2- → 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. O^2- + F₂

2. Al + O₂

3. K + Zn⁺²

Balance each half reaction in basic solution.

4. Cr₂O₇^{2 -} → Cr³⁺

5. NO → NO₃^-

6. SO₄^2- → SO₂

7. MnO₂ → Mn₂O₃

Balance each redox reaction in acid solution using the half reaction method.

8. H₂O₂ + Cr₂O₇^2- → O₂ + Cr³⁺

9. TeO₃^2- + N₂O₄ → Te + NO₃^-

10. ReO₄^- + IO^- → IO₃^- + Re

11. PbO₂ + I₂ → Pb²⁺ + IO₃^-

12. As → H₂AsO₄^- + AsH₃

Balance each redox reaction in basic solution using the half reaction method.

13. O₂ + Cr³⁺ → H₂O₂ + Cr₂O₇^2-

14. Te + NO₃^- → TeO₃^2- + N₂O₄

15. IO₃^- + Re → ReO₄^- + IO^-

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 → O^2-

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) F^1-

e) N₂

f) O^2-

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 O^2- 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- → S^2-

12. MnO₂ → MnO₄^-

13. Cr₂O₇^2- → CrO₄^2-

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.

SO₄^2- + Br₂ → S₂O₃^2- + 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. SO₄^2- + Br₂ → S₂O₃^2- + BrO₃^-

Classify as an oxidizing agent, reducing agent or both based on its position on the table.

State the E^o 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 E^o for each. Calculate the E^o 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? E^o = 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^- + 3SO₃^2- + 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 FeSO_4(aq)? Write a reaction and calculate E^o.

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^- + 3SO₃^2- + 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. ZnCl_2(l) electrolytic cell (electrowinning)

Anode: Cathode:

Anode reaction: Cathode reaction:

Overall reaction: MTV:

3. CuSO_4(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 FeI_2(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 CaCl_2(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 AgNO_3(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. SO₄^2-

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	U^3+	Y	reaction
2	V^2+	U	reaction
3	V^2+	Y	reaction
4	Y^3+	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) + I_2(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) + 3Cl_2(aq) + 7H₂O_(l) → Cr₂O₇^2-_(aq) + 6Cl^-_(aq) + 14H⁺_(aq)

The species which loses electrons is

A. Cl₂

B. Cr³⁺

C. H₂O

D. Cr₂O₇^2-

Quiz #3 Balancing Redox reactions- Acid & Base Cell Potentials

1. Consider the following overall reaction:

2Rh⁺ + Pb_(s) → 2Rh_(s) + Pb^2+- 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^- → 3SO₄^2- + 6H⁺ + Cl^-

The reduction half-reaction is

A. ClO₃^- + 6H⁺ → Cl^- + 3H₂O + 6e^-

B. ClO₃^- + 6H⁺ + 6e^- → Cl ^- + 3H₂O

C. SO₂ + 2H₂O → SO₄^2- + 4H⁺ + 2e^-

D. SO₂ + 2H₂O + 2e^- → SO₄^2- + 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 CrO₄^2-

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. H₂SO_3, H₃PO₄, H₂O₂

C. H₃PO₄, H₂SO_{3 ,} H₂O₂

D. H₂O_2,H₂SO_{3 ,} 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. SO₄^2- 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 E^o is positive.

B. spontaneous and E^o is negative.

C. non-spontaneous and E^o is positive.

D. non-spontaneous and E^o 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 H_2(g) and O_2(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 25^oC 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⁺ → Cr₂O₇^2-

D. CrO₄^2- → 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 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. 2SO₄^2- → S₂O₈^2- + 2e^-

C. H₂O → ½ O_2(g) + 2H⁺ + 2e^-

D. 2H₂O + 2e^- → H_2(g) + 2OH^-

 9. In the electrolysis of molten PbBr₂, the products at the anode and cathode are

A. B. C. D. ANODE (INERT)	CATHODE (INERT)
Br2	H2
O2	Pb
Pb	Br2
Br2	Pb

10. Under which conditions could an electrochemical cell provide 0.93V?

A. B. C. D. Anode	Cathode
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. SO₄^2-

D. Cr₂O₇^2-

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 E^o (volts) of the following half-reaction:

Pd²⁺ + 2e^- → Pd Pd²⁺ reacts with Cu_(s) but not with Hg_(l).

Based on the above, the E^o (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 E^o 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

Cu^2+

Fe^2+2

Zn^2+

Au^3+

2) Which most readily loses electrons?

Hg(l)

Cu^2+

Sn^4+

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 E^o

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 E^o.

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 E^o = +0.64 V

3V²⁺ + 2Al → 3V + 2Al³⁺ E^o = +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) → HV₂O₇^3-

16. Balance the following redox reaction occurring in basic solution:

MnO₄^- + C₂O₄^2- → MnO₂ + CO₂

17. 250.0 ml 0.200M MnO_4- reacts with excess SO₃^2-. How many grams of MnO₂ are produced?

2MnO₄^- + 3SO₃^2- +H₂O → 2MnO₂ +3SO₄^2- + 2OH^-

18. Determine the oxidation number for each bold atom.

MnO2	IO3^-	Cr2O7^2-	C2O4^2-	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 + 3SO₄^2-+ 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.

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	E^o
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 Cr2O7^2- 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⁺ + 2SO₄^2- → 2PbSO₄ + 2H₂O

The balanced, reduction half reaction is

A. Pb + SO₄^2- → 2PbSO₄ + 2e^-

B. Pb + 2H⁺ + SO₄^2- → PbSO₄ + 2H₂O + 2e^-

C. PbO₂ + 4H⁺ + SO₄^2- + 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. N_a 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 E^o 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 E^o is positive

B. spontaneous and E^o is negative

C. non-spontaneous and E^o is positive

D. non-spontaneous and E^o 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^- + 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.

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.

voltmeter

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⁺ + Te^2- → no reaction

U⁴⁺ + Te^2- → 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⁺ Te^2- V²⁺

D V²⁺ Te^2- 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 I_2(s)

B Ni_(s)

C Zn_(s)

D Br_2(l)

14. The oxidation number of carbon in C₂O₄^2- 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 25^oC, 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 E^o 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 + 2S₂O₃^2- → Ag(S₂O₃)₂^3- + 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⁺ + Te^2- → no reaction

U⁴⁺ + Te^2- → 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⁺ Te^2- V²⁺

D V²⁺ Te^2- 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 2SO₄^2- → S₂O₈^2- + 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:

NiO_2(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 O_2(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 + SO₄^2- + 2H⁺ → Se + H₂SO₃ + H₂O

Calculate the E^o 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