Oxidation and Reduction Questions

Power Point Lesson Notes- double click on the lesson number.

Worksheets Quiz

1. Oxidation, Reduction, Agents, & Reactions. WS 1

2. Lab: The Strength of Oxidizing Agents. Lab Handout

3. Oxidation Numbers Spontaneous Reactions WS 2

4. Oxidation Numbers, Application to Reactions. WS 3 1

5. Balancing Redox Half Reactions Acid/Base. WS 4

6. Balancing Redox Reactions in Acid/Base. WS 5 2

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. WS # 11 Practice Test # 1

15. Review WS # 12 Practice Test # 2

16. Review

17. Test.

Text book Hebden Read Unit V

If you want an A in this class you need to do this!!

Redox Half Reactions and Reactions WS #1

Define each

1. Oxidation - loss of electrons

2. Reduction - gain of electrons

3. Oxidizing agent - causes oxidation by undergoing reduction

4. Reducing agent - causes reduction by undergoing oxidation

Write half reactions for each of the following atoms or ions. Label each as oxidation or reduction.

5. Al -----------> Al³⁺ + 3e^- oxidation

6. S + 2e^- ---------> S^2- reduction

7. 2O^2- ----------> O₂ + 4e^- oxidation

8. Ba²⁺ + 2e^- -----------> Ba reduction

9. 2N^3- ----------> N₂ + 6e^- oxidation

10. Br₂ + 2e^- ---------> 2Br^- reduction

11. P + 3e^- ----------> P^3- reduction

12. Ca -----------> Ca²⁺ + 2e^- oxidation

13 Ga³⁺ + 3e^- -----------> Ga reduction

14. S + 2e^- ---------> S^2- reduction

15. H₂ ---------> 2H⁺ + 2e^- oxidation

16. 2H⁺ + 2e^- ---------> H₂ reduction

17. 2F^- ----------> F₂ + 2e^- oxidation

18. P^3- ----------> P + 3e^- oxidation

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

19. Al & Zn²⁺

2Al + 3Zn²⁺ → 2Al³⁺ + 3Zn

oxidized reduced

reducing agent oxidizing agent

20. F₂ & O^2-

2F₂ + 2O^2- → 4F^- + O₂

reduced oxidized

oxidizing agent reducing agent

21. O₂ & Ca

2Ca + O₂ → 2Ca²⁺ + 2O^2-

oxidized reduced

reducing agent oxidizing agent

22. Al³⁺ & Li

Al³⁺ + 3Li → Al + 3Li⁺

reduced oxidized

oxidizing agent reducing agent

Label the species that is reduced, that is oxidized, the reducing agent and the oxidizing agent.

23. Fe²⁺ + Co → Co²⁺ + Fe

Co → Co²⁺ + 2e^- oxidation Fe²⁺ + 2e- → Fe reduction

24. 3 Ag⁺ + Ni → Ni³⁺ + 3 Ag

Ni → Ni³⁺ + 3e^- oxidation Ag⁺ + 1e- → Ag reduction

25. Cu²⁺ + Pb → Pb²⁺ + Cu

Pb → Pb²⁺ + 2e^- oxidation Cu²⁺ + 2e- → Cu reduction

26. O₂ + 2 Sn → O^2- + 2 Sn²⁺

Sn → Sn²⁺ + 2e^- oxidation O₂ + 4e- → 2O^2- reduction

27. Co²⁺ + 2 F^- → Co + F₂

2F^- → F₂ + 2e^- oxidation Co²⁺ + 2e- → Co reduction

28. List the species (formulas from above) that lose electrons:

Co Ni Pb Sn F^-

29. List the species (formulas from above) that gain electrons:

Fe²⁺ Ag⁺ Cu²⁺ O₂ Co²⁺

For each of the following reactions, identify:

-The Oxidizing Agent.

-The Reducing Agent.

-The Substance Oxidized.

-The Substance Reduced.

30. I^- + Cl₂ ----------> Cl^- + I₂

Substance oxidized I^- Reducing agent I^-

Oxidizing agent Cl₂ Substance reduced Cl₂

31. Co + Fe³⁺ -----------> Co²⁺ + Fe²⁺

Substance oxidized Co Reducing agent Co

Oxidizing agent Fe³⁺ Substance reduced Fe³⁺

32. Cr⁶⁺ + Fe²⁺ -----------> Cr³⁺ + Fe³⁺

Substance oxidized Fe²⁺ Reducing agent Fe²⁺

Oxidizing agent Cr⁶⁺ Substance reduced Cr⁶⁺

Redox Half Reactions and Reactions WS #2

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

a) NH₃ -3 b) H₂SO₄ 6

c) ZnSO₃ 4 d) Al(OH)₃ 3

e) Na 0 f) Cl₂ 0

g) AgNO₃ 5 h) ClO₄^- 7

i) SO₂ 4 j) K₂Cr₂O₄ 3

k) Ca(ClO₃)₂ 5 l) K₂Cr₂O₇ 6

m) HPO₃^2- 3 n) HClO 1

o) MnO₂ 4 p) KClO₃ 5

q) PbO₂ 4 r) PbSO₄ 2

s) K₂SO₄ 6 t) NH₄⁺ -3

u) Na₂O₂ -1 v) FeO 2

w) Fe₂O₃ 3 x) SiO₄^4- -2

y) NaIO₃ 5 z) ClO₃^- 5

aa) NO₃^- 5 bb) Cr(OH)₄ 4

cc) CaH₂ -1 dd) Pt(H₂0)₅(0H)²⁺ +3

ee) Fe(H₂O)₆³⁺ +3 ff) CH₃COOH 0

2. What is the oxidation number of carbon in each of the following substances?

a) CO 2 b) C 0

c) CO₂ 4 d) CO₃^2- 4

e) C₂H₆ -3 f) CH₃OH -2

3. For each of the following reactions, identify: the oxidizing agent, the reducing agent, the substance oxidized and the substance reduced.

a) Cu²⁺ _(aq) + Zn _(s) --------> Cu _(s) + Zn²⁺ _(aq)

Substance oxidized Zn Substance reduced Cu²⁺

Oxidizing agent Cu²⁺ Reducing agent Zn

b) Cl_{2 (g)} + 2 Na _(s) --------> 2 Na+ _(aq) + 2 Cl^- _(aq)

Substance oxidized Na Substance reduced Cl₂

Oxidizing agent Cl₂ Reducing agent Na

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

            O₂ Cl^- Fe                                Na⁺

5.         Determine the oxidation number for the element underlined.

            PbSO₄ +6 ClO₃^- +5

            HPO₃^2- +3                              Na₂O₂ -1

            CaH₂ -1                                 Al₂(SO₄)₃          +6

            NaIO₃ +5 C₄H₁₂ -3

6.         Al³⁺    +       Zn                        →        Al        +          Zn²⁺

+3                  0                                     0                         +2       oxidation numbers!    

            Substance oxidized    Zn           Oxidizing agent    Al³⁺

7.         Cr₂O₇^2-            +          ClO₂^-    →         Cr³⁺     +          ClO₄^-

+6                               +3                  +3                    +7            oxidation numbers  

            Substance reduced    Cr₂O₇^2-           Oxidizing agent Cr₂O₇^2-

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

            a) NH₃ -3                               b) H₂SO₄ +6

            c) ZnCO₃             +4                                d) Al(OH)₃ +3

            e) Na 0                                  f) Cl₂ 0

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

                        Al        &         AgNO₃

Al     →  Al³⁺    +   3e^-

                        3(Ag⁺   +  1e^-    →     Ag)

            Al        +          3Ag⁺   →     Al³⁺     +     3Ag

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

            O₂ Cl^- Fe                    Na⁺

11.       Cr³⁺       +        ClO₄^-                       →        Cr₂O₇^2- +          ClO₂^-

+3                      +7                                       +6                  +3            oxidation numbers 

            Substance reduced     ClO₄^-     Oxidizing agent         ClO₄^-

12.       O₃   +    H₂O   +  SO₂  →    SO₄^2-    +    O₂     +   2H⁺

?                        +4            +6                0            oxidation numbers

                 Substance oxidized               SO₂   Reducing  agent      SO₂

13.       3As₂O₃    +    4NO₃^-  +  7H₂O   +  4 H⁺   →  6H₃AsO₄    +   4NO

+3                 +5                                                  +5               +2            oxidation numbers

                 Substance reduced     NO₃^-                            Reducing agent            As₂O₃

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

HP0₃^2- 3 Na₂O₂ -1

CaH₂ -1 Al₂(SO₄)₃ 6

NaIO₃ 5 C₄H₁₂ -3

14. Al³⁺ + Zn ---------> Al + Zn²⁺

Substance oxidized Zn Oxidizing agent Al³⁺

15. Cr₂O₇^2- + ClO₂^{- ------------>} Cr³⁺ + ClO₄^-

Substance reduced Cr₂O₇^2- Oxidizing agent Cr₂O₇^2-

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

a) NH₃ -3 b) H₂SO₄ 6

c) ZnCO₃ 4 d) Al(OH)₃ 3

e) Na 0 f) Cl₂ 0

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

Al + 3Ag⁺ ----------> Al³⁺ + 3Ag

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

O₂ Cl^- Fe Na⁺

Determine the oxidation number for the element underlined.

19. PbSO₄   2

20. ClO₃^- 5

21. HPO₃^2- 3

22. Na₂0₂ -1

23. CaH₂ -1

24. NaIO₃ 5

25. C₄H₁₂ -3

26. Al₂(SO₄)₃ 6

27. Al³⁺ + Zn ----------> Al + Zn²⁺

Substance oxidized Zn Oxidizing agent Al³⁺

28. Cr₂O₇^2- + ClO₂^- ----------------> Cr³⁺ + ClO₄^-

Substance reduced Cr₂O₇^2- Oxidizing agent Cr₂O₇^2-

29. O₃ + H₂O + SO₂ -----> SO₄^2- + 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- --------------> 2SO₄^2- + 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^- + 5e^- -------------> N + 3H₂O

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. 2H₂O + 2IVO₃ + 8e^- --------------> I₂ + 2VO₂^2- + 4OH^-

13. 2IVO₃ + 2e^- --------------> I₂ + 2VO₃^-

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

15. 4H₂O + 6e^- + SO₄^2- --------------> S + 8OH^-

16. 10 OH^- + S₂O₃^2- --------------> 2SO₄^2- + 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.

SO₃^2- --------> SO₄^2- oxidation

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

CrO₄^2- --------> Cr₂O₇^2- neither

CrO₄^2- --------> Cr³⁺ reduction

2I^- --------> I₂ oxidation

IO₃^- --------> I₂ reduction

MnO₄^- --------> Mn²⁺ reduction

ClO₂^- --------> ClO^- reduction

20. Cr₂O₇^2- + Fe²⁺ --------> Cr³⁺ + Fe³⁺

Substance oxidized Fe²⁺ Substance reduced Cr₂O₇^2-

Oxidizing agent Cr₂O₇^2- 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. 2O^2- + 2F₂ -----------> O₂ + 4F^-

2. 4Al + 3O₂ -----------> 6O^2- + 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- + SO₄^{2- -------------->} SO₂ + 4OH^-

7. 2MnO₂ + H₂O + 2e- ^{-------------->} Mn₂O₃ + 2OH^-

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

8. 8H⁺ + 3H₂O₂ + Cr₂O₇^2- -------> 3O₂ + 2Cr³⁺ + 7H₂O

9. TeO₃^{2 - +} 2N₂O₄ + H₂O -------> Te + 4NO₃^- + 2H⁺

10. 12H₂O +          10V            ------->        6H⁺       +     3H₂V₂O₄^2-   +          4VH₃

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. 10IVO₃ + 28OH^- -------> 10VO₂^2- + I₂ + 8IO₃^- 14H₂O

16. 8OH^- + 5Pb²⁺ + 2IO₃^- -------> 5PbO₂ + I₂ + 4H₂O

17.       2H₂O      +  O₂ +    4Fe(H₂O)₆²⁺            ------->  4Fe(H₂O)₆³⁺    +      4OH^- 

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 ^--------> O^2- reduction

22. P₂O₅ ^--------> P₄H₁₀ reduction

Determine the oxidation number

23. H₂SO₄ 6 22. HSO₄^- 6

24. P₄ 0 23. NaH -1

25. UO₃ 6 24. Na₂O₂ -1

26. U₂O₅ 5 25. PbSO₄ 2

WS #6 Review

1. Describe each in your own words

1. Oxidation - loss of electrons

2. Reduction - gain of electrons

3. Oxidizing agent - causes oxidation by undergoing reduction

4. Reducing agent - causes reduction by undergoing oxidation

2. Write half reactions for each. Describe as oxidation or reduction. Circle all oxidizing agents.

a) Na -----------> Na⁺ + e^- oxidation

b) Ca -----------> Ca²⁺ + 2e^- oxidation

c) Al³⁺ + 3e^- -----------> Al reduction

d) 2F^1- ----------> F₂ + 2e^-           oxidation

e) N₂ + 6e^- ----------> 2N^3- reduction

f) 2O^2- ----------> 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 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.

strongest oxidizing agent Ni²⁺ + 2e^- -----------> Ni

Mn²⁺ + 2e^- -----------> Mn

Al³⁺ + 3e^- -----------> Al strongest reducing agent

7. Ag⁺ reacts with Pb, however, Ca⁺² does not react with Pb. Rank the reducing agents in order of decreasing strength. Rank the oxidizing agents in order of decreasing strength.

strongest oxidizing agent Ag⁺ + 1e^- -----------> Ag

Pb²⁺ + 2e^- -----------> Pb

Ca²⁺ + 2e^- -----------> Ca strongest reducing agent

8. Cl₂ reacts with Ag, however, Ag does not react with Mg⁺². Rank the oxidizing agents in order of decreasing strength. Rank the reducing agents in order of decreasing strength.

strongest oxidizing agent Cl₂ + 2e^- --------> 2Cl^-

Ag⁺ + 1e^- ----------->    Ag

Mg²⁺ + 2e^- -----------> Mg strongest reducing agent

9. Ni⁺² reacts with Mn, however, Al⁺³ does not react with Mn. Rank the reducing agents in order of decreasing strength. Rank the oxidizing agents in order of decreasing strength.

strongest oxidizing agent Ni²⁺ + 2e^- -----------> Ni

Mn²⁺ + 2e^- -----------> Mn

Al³⁺ + 3e^- -----------> Al strongest reducing agent

10. Cl₂ reacts with Br^-, however, I₂ does not react with Br^-. Rank the oxidizing agents in order of decreasing strength. Rank the reducing agents in order of decreasing strength.

strongest oxidizing agent Cl₂ + 2e^- --------> 2Cl^-

Br₂ + 2e^- --------> 2Br^-

I₂ + 2e^- --------> 2I^- strongest reducing agent

Classify as oxidation, reduction or neither.

11. SO₄^2- --------> S^2- reduction

12. MnO₂ --------> MnO₄^- oxidation

13. Cr₂O₇^2- --------> CrO₄^2- neither

14. IO₃^- --------> I₂ reduction

15. Given the following lab data

SnCl₂ & Ni Spontaneous

Ni(NO₃)₂ & Fe Spontaneous

Cr(NO₃)₃ & Fe Non spontaneous.

i) Write three balanced equations.

Ni + Sn²⁺ -------------> Ni²⁺ + Sn

Fe + Ni²⁺ -------------> Fe²⁺ + Ni

Fe + Cr³⁺ <------------- Fe²⁺ + Cr

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

strongest oxidizing agent Sn²⁺ + 2e^- -----------> Sn

Ni²⁺ + 2e^- -----------> Ni

Fe²⁺ + 2e^- -----------> Fe

Cr³⁺ + 3e^- -----------> Cr strongest reducing agent

iii) Rank the reducing agents in decreasing order of strength. See above.

iv) Will SnCl₂ react with Cr? Explain? Yes, because Sn²⁺ is a stronger oxidizing agent than Cr^{3+ .}

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 ----------> 5S₂O₃^2- + 8BrO₃^- + H₂O

oxidizing agent reducing agent

18.      Balance the redox equation in acid solution    IPO₄  →   I₂   +    IO₃^-    +    PO₄^3-

           Hint: You need two half reaction

IPO₄   →   I₂                 +        PO₄^3-

          IPO₄   →  IO₃^-     +      PO₄^3-

9H₂O   +   5IPO₄    →    3IO₃^-    +    I₂   +    5PO₄^3-    +   18H⁺

19. Describe as spontaneous or non-spontaneous. Use your reduction potential chart.

a) ZnCl₂ & Cu      nonspontaneous

b) CuCl₂ & NaCl nonspontaneous

c) Br₂ & Fe²⁺ spontaneous

d) H₂S & Al³⁺ nonspontaneous

20. Can you keep HCl in a Zn container? No, Spontaneous reaction.

What about an Au container? Yes, nonspontaneous reaction.

Balance in basic solution

21. H₂O + 10SO₄^2- + 4Br₂ ------> 5S₂O₃^{2- +} 2OH^- + 8BrO₃^-

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

State the E^oor voltage of its position. Some of these are both, so state two voltages and indicate that it can be an oxidizing and reducing agent.

e.g. MnO₄^- (in acid) oxidizing agent 1.51 v

22. Br₂ oxidizing agent     1.09 v

23. Fe²⁺ oxidizing agent / reducing agent -0.45 v / - 0.77 v

24. MnO₄^- (water) oxidizing agent 0.60 v

25. Ni reducing agent 0.26 v

26. Cr³⁺ oxidizing agent -0.74 v / -0.41v

27. H₂O oxidizing agent / reducing agent     -0.41 v / -0.82 v

Indicate as spontaneous or non-spontaneous.

28. MnO₄^- & Fe²⁺ non-spontaneous

29. Cu²⁺ & Br^- non-spontaneous

30. HNO₃ & Ag spontaneous

31. MnO₄^- (acid) & H₂O spontaneous

32. Ni_(s) & Al³⁺ non-spontaneous

33. HCl & Mg spontaneous

Write each oxidation and reduction half reaction for each question above. Determine the E^o for each. Calculate the E^o for the overall reaction.

34. MnO₄^- + 2H₂O + 3e^- --------> MnO₂ + 4OH^- +0.60 v

3(Fe²⁺ -----------> Fe³⁺ + 1e^-) -0.77 v

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

35.

36. NO₃^- + 4H⁺ +3e^- -----------> NO + 2H₂O +0.96 v

3(Ag ----------> Ag⁺ + 1e^-) -0.80 v

NO₃^- + 4H⁺ + 3Ag ----------> NO + 2H₂O + 3Ag+ +0.16 v

37.

38.

39. 2H⁺ + 2e^- ------> H₂ 0.00 v

Mg ----------> Mg²⁺ + 2e^- 2.37 v

Mg + 2H⁺ ----------> Mg²⁺ + H₂ 2.37 v

WS # 7 Electrochemical Cells

1. Oxidation is when electrons are lost.

2. Reduction is when electrons are gained.

3. The reducing agent undergoes oxidation.

4. The oxidizing agent undergoes reduction.

5. A negative voltage means the reaction is nonspontaneous.

6. In an electrochemical cell electrons exit the electrode, which is negative.

7. In an electrochemical cell the reduction reaction is higher on the chart, while the

oxidation reaction is lower. .

8. The cathode is the site of reduction and the anode is the site of oxidation. .

9. Anions migrate to the anode and cations migrate to the cathode.

10. Anions have a negative charge and cations have a positive charge.

Draw and completely analyze each electrochemical cell.

11. Zn / Zn(NO₃)₂ ║ Cu / Cu(NO₃)₂

12. Ag / AgNO₃ ║ H₂ / HCl

WS # 8

1. In an electrolytic cell, reduction occurs at the negative electrode and oxidation occurs at the positive electrode.

2. If there are two possible reduction reactions, the highest one on the chart occurs.

3. For reduction, the chart is read from left to right.

4. For oxidation, the chart is read from right to left and the sign of the voltage is changed.

5. If there are two possible oxidation reactions, the lowest one on the chart occurs.

6. Corrosion of a metal is oxidation.

7. Electrolysis uses electrical energy.

8. Electrochemical cells produce electrical energy.

9. Electrolytic cells use electrical energy.

10. What is the standard reference cell? hydrogen E^o = O v

Draw and completely analyze each electrolytic cell.

11. Molten NaCl

Cathode: Na⁺ + 1e^- → Na_(s) -2.71 v Anode: 2Cl^- → Cl₂ + 2e^- -1.36 v

Overall: 2Na⁺ + 2Cl^- → Cl₂ + 2Na_(s) -4.07 v MTV = +4.07 v

12. Aqueous Na₂SO₄

Cathode: 2H₂O + 2e^- → H₂ + 2OH^- -0.41 v Anode: H₂O → 2H⁺ + 1/2O₂ + 2e^- -0.82 v

Overall: H₂O → H₂ + 1/2O₂ -1.23 v MTV = +1.23 v

13. Liquid K₂O

Cathode: K⁺ + 1e^- → K_(s) -2.93 v Anode: 2O^2- → O₂ + 4e^- ? v

Overall: 4K⁺ + 2O^2- → O₂ + 4K_(s) -? v MTV = +? v

14. 1.0 M LiI

Cathode: Cathode: 2H₂O + 2e^- → H₂ + 2OH^- -0.41 v Anode: 2I^- → I₂ + 2e^- -0.54 v

Overall: 2H₂O + 2I^- → I₂ + H₂ + 2OH^- -0.95 v MTV = +0.95 v

15. 250ml of 0.200M MnO₄^- reacts with excess SO₃^-2. How many grams of MnO₂ are produced? This is Chemistry 11 stoichiometry. 2MnO₄^- + 3SO₃^-2 + H₂O -----> 2MnO₂ + 3SO₄^-2 + 2OH^-

0.250L MnO₄^- x 0.200 mol x 2 mol MnO₂ x 86.9g = 4.34g

L 2 mol MnO₄^- mol

16. Determine the oxidation number for each underlined atom.

MnO₂ 4 Cr₂O₇^-2 6 IO₃^- 5 C₂O₄^-2 3 Al(NO₃)₃ 5

17. Describe each term:

Salt bridge- a u-tube filled with salt solution that allows ions to flow in an electrochemical cell.

Electrolyte- a solution that conducts electricity

Anode- an electrode that is the site of oxidation

Cathode- an electrode that is the site of reduction

Spontaneous- a reaction that occurs naturally and has a positive voltage

Electron affinity- the ability of a metal to attract electrons

18. What would happen if you used an aluminum spoon to stir a solution of FeSO_4(aq) ? Write a reaction and calculate E^o.

2Al + 3Fe²⁺ -------> 2Al³⁺ + 3Fe E⁰ = 1.21 v Spontaneous. There would be a reaction!

19. Draw an electrochemical cell using Cu and Ag electrodes.

Cathode (+) Anode (-)

Ag Cu

Ag⁺ + 1e^---------> Ag 0.80v Cu -------> Cu² + 2e -0.34v

2Ag⁺ + Cu ------> 2Ag + Cu²⁺ E⁰ = 0.46 v spontaneous

20. 250ml of .500M MnO₄^- are required to titrate a 100ml sample of SO₃^-2. Calculate the [SO₃^-2]

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

.250L MnO₄^- x 0.500 mol x 3 mol SO₃^-2

L 2MnO₄^- = 1.88M

0.100L

21. How is the breathalyzer reaction used to determine blood alcohol content (you might need to look this up in your textbook)?

The breathalyzer reaction uses a spontaneous redox reaction between acidic 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. ZnCl_2(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. CuSO_4(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: 2O^2- -------> O₂ + 4e^- Cathode reaction: Al³⁺ + 3e^- -------> Al

Overall reaction: 6O^2- + 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: 2O^2- -------> O₂ + 4e^- Cathode reaction: Al³⁺ + 3e^- -------> Al

Overall reaction: 6O^2- + 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 FeI_2(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. 2HIO₃ + 5H₂SO₃ ----------> I₂ + 5H₂SO₄ + H₂O

oxidizing agent HIO₃ substance oxidized H₂SO₃

substance reduced HIO₃ reducing agent H₂SO₃

12. What is the fuel in a fuel cell? H₂ and O₂

13. Describe the differences and similarities between an electrolytic and electrochemical cell.

Electrolytic Electrochemical

Uses electricity Produces electricity

Nonspontaneous Spontaneous

Makes chemicals Uses chemicals

Inert carbon electrodes Usually has a salt bridge

The 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.

14. 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.

15. Describe and give one example of electrorefinning. The electrorefinning of Pb.

16. List three metals that can be won from aqueous solution. Pb Au Ag Zn Cu Fe Sn

17. List three metals that cannot be won from aqueous solution. Na K Li Ca Mg Al

18. 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.

19. Write the half reaction that describes the corrosion of iron. Fe --------> Fe²⁺ +2e^-

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

21. 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.

22. Write the anode and cathode reaction in an electrolytic cell with a CaCl₂ (l) electrolyte.

Cathode: Ca²⁺ + 2e^- ---------> Ca Anode: 2Cl^- ----------> Cl₂ + 2e^-

23. 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.

24. 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.

25. Describe as an electrochemical or electrolytic cell:

a) Fuel cell electrochemical b)Charging a car battery electrolytic

c) Discharging a car battery electrochemical d) Ni plating electrolytic

e) Industrial Al production electrolytic f) Cl₂ production electrolytic

26) 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

29.       Analyze This

Label each anode and cathode. See Diagram
Write each anode and cathode reaction. See Diagram
Indicate the ion migration in each cell. See Diagram
Determine the initial cell voltage of the electrochemical cell. See Diagram
Determine the MTV for the electrolytic cell. See Diagram
Will electrolysis occur? Yes
Indicate electron flow. See Diagram
Indicate all electrodes that gain mass. Ag and Cu
Indicate all electrodes that lose mass. Mg
What happens to [NO3-] in the Mg half-cell? Increases
What happens to the [Ag+] in the Ag half-cell? Decreases
What happens to [Mg2+] in the Mg half-cell? Increases
What is the equilibrium electrochemical cell potential? 0 V
What chemical is made at the Pt electrode on the right? Cu
What chemicals are made at the Pt electrode on the left? O₂ and H⁺