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Oxidation and Reduction Reactions Workbook Period/Topic Worksheets 1. Oxidation, Reduction, Agents, & Reactions. 2. Lab: The Strength of Oxidizing Agents. 3. Oxidation Numbers Spontaneous Reactions WS 2 4. Oxidation Numbers, Application to Reactions. WS 3 5. Balancing Redox Half Reactions Acid/Base. WS 4 6. Balancing Redox Reactions in Acid/Base. WS 5 7. Standard Potentials Using Chart. WS 6 8. Electrochemical Cells. WS 7 9. Electrochemical Cells Lab. 10. Electrolytic Cells. 11. Electrolytic Cells Lab. 12. Application of Electrochemical Cells 13. Quiz WS 1 1 2 3 WS 8 4 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. 2. 3. 4. Oxidation Reduction Oxidizing agent Reducing agent Write half reactions for each of the following atoms or ions. Label each as oxidation or reduction. 5. Al 6. S 7. O-2 8. Ba2+ 9. N3- 10. Br2 11. P 12. Ca 13 Ga3+ 14. S 15. H2 16. H+ 17. F- 18. P3- Balance each spontaneous redox equation. Identify the entities reduced and oxidized. State the reducing agent and the oxidizing agent. 19. Al & Zn2+ 20. F2 & O2- 21. O2 & Ca 22. Al3+ & Li Write the oxidation and reduction reactions for each redox reaction. The first one is done for you. 23. 24. Fe2+ + ⇄ Co Oxidation: Co Co2+ + 2e- Reduction: Fe2+ + 2e- Fe 3 Ag+ Co2+ + Fe + Ni ⇄ Ni3+ + 3 Ag + Pb ⇄ Pb2+ + Cu + 2 Sn ⇄ O2- + 2 Sn2+ Oxidation: Reduction: 25. Cu2+ Oxidation: Reduction: 26. O2 Oxidation: Reduction: Co2+ 27. + 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 1. 2. Redox Half Reactions and Reactions State the Oxidation Number of each of the elements that is underlined. a) NH3 _____ b) H2SO4 _____ c) ZnSO3 _____ d) Al(OH)3 _____ e) Na _____ f) Cl2 _____ g) AgNO3 _____ h) ClO4- _____ i) SO2 _____ j) K2Cr2O4 _____ k) Ca(ClO3)2 _____ l) K2Cr2O7 _____ m) HPO32- _____ n) HClO _____ o) MnO2 _____ p) KClO3 _____ q) PbO2 _____ r) PbSO4 _____ s) K2SO4 _____ t) NH4+ _____ u) Na2O2 _____ v) FeO _____ w) Fe2O3 _____ x) SiO44- _____ y) NaIO3 _____ z) ClO3- _____ aa) NO3- _____ bb) Cr(OH)4 _____ cc) CaH2 _____ dd) Pt(H2O)5(OH)2+ _____ ee) Fe(H2O)63+ _____ ff) CH3COOH _____ What is the oxidation number of carbon in each of the following substances? a) CO _____ b) C _____ c) CO2 _____ d) CO32- _____ e) C2H6 _____ f) CH3OH _____ 3. For each of the following reactants, identify: the oxidizing agent, the reducing agent, the substance oxidized and the substance reduced. a) Cu2+ (aq) b) + → Zn (s) Substance oxidized Oxidizing agent _____ _____ Cl2 (g) + → 2 Na (s) Substance oxidized Oxidizing agent Worksheet # 3 Cu(s) + Zn2+ (aq) Substance reduced Reducing agent 2 Na+ (aq) _____ _____ + _____ _____ 2 Cl- (aq) Substance reduced Reducing agent ____ _____ Spontaneous and Non-spontaneous Redox Reactions Describe each reaction as spontaneous or non-spontaneous. 1. Au3+ → Fe2+ + Au 2. Pb + → Fe2+ + Pb2+ 3. Cl2 + → F2 4. S2O82- + Pb → 2SO42- 5. Cu2+ + 2Br- → Cu 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. + Fe3+ Fe3+ F- + 2Cl+ Pb2+ + Br2 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 Cl- O2 13. Determine the oxidation number for the element underlined. PbSO4 HP032CaH2 NaIO3 14. Na+ Fe Al3+ __________ __________ __________ __________ + → Zn ClO3Na2O2 Al2(SO4)3 C4H12 __________ __________ __________ __________ Al Zn2+ + Substance oxidized _______ Oxidizing agent ________ 15. Cr2O72- + ClO2- → Cr3+ + ClO4- Substance reduced ________ Oxidizing agent ________ 16. State the Oxidation Number of each of the elements that is underlined. a) NH3 c) ZnCO3 e) Na __________ __________ __________ b) H2SO4 d) Al(OH)3 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. 20. 21. 22. 23. 24. 25. 26. 27. PbSO4 __________ - ClO3 HPO32Na202 CaH2 NaIO3 C4H12 Al2(SO4)3 Al3+ + Zn Substance oxidized 28. Cr2O72- + + __________ Oxidizing agent ClO2- → Substance reduced 29. __________ __________ __________ __________ __________ __________ __________ → Al Cr3+ __________ 3As2O3 __________ + 2H+ Reducing agent + 4NO3- + 7H2O + 4 H+ → 6H3AsO4 Substance reduced __________ __________ ClO4- Oxidizing agent O3 + H2O + SO2 → SO42- + O2 Substance oxidized__________ 30. + Zn2+ Reducing agent __________ + 4NO __________ Worksheet # 4 Balancing Redox Reactions Balance each of the following half-cell reactions. (In each case assume that the reaction takes place in an ACIDIC solution.) Also, state whether the reaction is oxidation or reduction. 1. S2O32- SO42- 2. MnO4- 3. As → AsO43- 4. Cr3+ → Cr2O72- 5. Pb2+ → PbO2 6. SO42- → S 7. NO3- → NO 8. NO3- → NH4+ 9. BrO3- → Br2 → Mn2+ → Balancing Half Cell Reactions Balance in basic solution. 10. NO3- → NO 11. MnO4- → Mn2+ 12. As → AsO43- 13. Cr3+ → Cr2O72- 14. Pb2+ → PbO2 15. SO42- → S 16. S2O32- → SO42- 17. NO3- → NH4+ 18. BrO3- → Br2 19. Determine if each of the following changes is oxidation, reduction or neither. SO32- → SO42- ________________ CaO → Ca ________________ CrO42- → Cr2O72- ________________ CrO42- → Cr3+ ________________ 20. 2I- → I2 ________________ IO3- → I2 ________________ MnO4- → Mn2+ ________________ ClO2- → ClO- ________________ Cr2O72- + Fe2+ Substance oxidized Oxidizing agent Worksheet # 5 → Cr3+ _____ _____ + Fe3+ Substance reduced Reducing agent _____ _____ Balancing Redox Reactions in Acid and Basic Solution Balance each redox equation. Assume all are spontaneous. Use the half reaction method. 1. O2- + F2 2. Al + O2 3. K + Zn+2 Balance each half reaction in basic solution. 4. Cr2O72 - → Cr3+ 5. NO → NO3- 6. SO42- → SO2 7. MnO2 → Mn2O3 Balance each redox reaction in acid solution using the half reaction method. 8. H2O2 + Cr2O72- → 9. TeO32- + N2O4 → 10. ReO4- + IO- → 11. PbO2 + I2 → 12. → As O2 Te + + IO3- NO3- + Pb2+ + IO3- H2AsO4- + AsH3 Balance each redox reaction in basic solution using the half reaction method. 13. O2 + Cr3+ 14. Te + NO3- → → H2O2 + Cr2O72- TeO32- + N2O4 Cr3+ Re 15. IO3- + Re → 16. Pb2+ + IO3- → 17. Cr2O72- + Hg → ReO4- + PbO2 Hg2+ IO- + + I2 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 21. O2 → O2- 22. P2O5 → P4H10 + H2O 22. HS O4- Determine the oxidation number 23. H2S O4 24. P4 23. NaH 25. U O3 24. Na2O 2 26. U 2O5 25. P b SO4 Worksheet # 6 1. Review 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) b) c) d) e) f) Na Ca Al3+ F1N2 O2- 3. Write the reaction between the following: Use the half reaction method. a) Ca + Al(NO3)3 b) Sn + AgNO3 c) Sn + Au(NO3)3 Cu Cu+ Al Al3+ F O2- O2 4. Circle each reducing agent: 5. Circle each oxidizing agent: F- 6. Ni+2 reacts with Mn, however, Al+3 does not react with Mn. Rank the oxidizing agents in order of decreasing strength. Rank the reducing agents in order of decreasing strength. 7. Ag+ reacts with Pb, however, Ca+2 does not react with Pb. Rank the reducing agents in order of decreasing strength. Rank the oxidizing agents in order of decreasing strength. 8. Cl2 reacts with Ag, however, Ag does not react with Mg+2. Rank the oxidizing agents in order of decreasing strength. Rank the reducing agents in order of decreasing strength. 9. Ni+2 reacts with Mn, however, Al+3 does not react with Mn. Rank the reducing agents in order of decreasing strength. Rank the oxidizing agents in order of decreasing strength. 10. Cl2 reacts with Br-, however, I2 does not react with Br-. Rank the oxidizing agents in order of decreasing strength. Rank the reducing agents in order of decreasing strength. Classify as oxidation, reduction or neither. 11. SO42- → S2- 12. MnO2 → MnO4- 13. Cr2O72- → CrO42- 14. IO3- → I2 15. Given the following lab data SnCl2 Ni(NO3)2 Cr(NO3)3 16. & & & Ni Fe Fe Spontaneous Spontaneous 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? Determine the oxidizing and reducing agent. Balance in acidic solution. MnO4- + H2S → S + MnO 17. Determine the oxidizing and reducing agent. Balance in acidic solution. SO42- + Br2 → S2O32- + MnO4- Balance in basic solution 19. Describe as spontaneous or non-spontaneous. Use your reduction potential chart. 20. ZnCl2 & Br2 Cu & 2+ Fe H2S → 18. a) c) + BrO3- b) d) S CuCl2 & H2S + MnO NaCl & Al3+ Can you keep HCl in a Zn container? Explain? What about an Au container? Balance in basic solution 21. SO42- + Br2 → S2O32- + BrO3- Classify as an oxidizing agent, reducing agent or both based on its position on the table. State the Eo or voltage of its position. Some of these are both, so state two voltages and indicate that it can be an oxidizing and reducing agent. e.g. MnO422. 23. 24. 25. 26. (in acid) Br2 Fe2+ MnO4- (water) Ni Cr3+ oxidizing agent _________________ _________________ _________________ _________________ _________________ 1.51 V _________________ _________________ _________________ _________________ _________________ 27. H2O _________________ _________________ Indicate as spontaneous or non-spontaneous. 28. 29. 30. MnO4- (Alkaline) & Fe2+ HNO3 & Ag HCl & Mg Write each oxidation and reduction half reaction for each question above. Determine the Eo for each. Calculate the Eo for the overall reaction. 34. 35. 36. Worksheet # 7 Electrochemical Cells 1. Oxidation is when electrons are . 2. Reduction is when electrons are . 3. The reducing agent undergoes . 4. The oxidizing agent undergoes . 5. A negative voltage means the reaction is 6. In an electrochemical cell electrons exit the electrode which is 7. In an electrochemical cell the reduction reaction is oxidation reaction is . 8. The cathode is the site of 9. Anions migrate to the 10. Anions have a . Zn / Zn(NO3)2 and cations migrate to the charge and cations have a ll on the chart, while the and the anode is the site of Draw and completely analyze each electrochemical cell. 11. . Cu / Cu(NO3)2 charge. . . 12. Ag / AgNO3 ll H2 / HCl Worksheet # 8 Electrolytic Cells 1. In an electrolytic cell, reduction occurs at the the electrode. 2. If there are two possible reduction reactions, the 3. For reduction, the chart is read from to 4. For oxidation, the chart is read from voltage is . to 5. If there are two possible oxidation reactions, the 6. Corrosion of a metal is 7. Electrolysis 8. Electrochemical cells 9. Electrolytic cells 10. What is the standard reference cell? Molten NaCl one on the chart occurs. . and the sign of the one on the chart occurs. . electrical energy. electrical energy. electrical energy. Draw and completely analyze each electrolytic cell. 11. electrode and oxidation occurs at Eo = v 12. Aqueous Na2SO4 13. Liquid K2O 14. 1.0 M LiI 15. 250.0 mL of 0.200 M MnO4- reacts with excess SO3-2. How many grams of MnO2 are produced? This is Chemistry 11 stoichiometry. 2MnO4- + 3SO32- + H2O → 2MnO2 + 3SO42- + 2OH- 16. Determine the oxidation number for each underlined atom. MnO2 Cr2O72IO3C2O42Al(NO3)3 17. Describe each term: Salt bridge Electrolyte Anode Cathode Spontaneous Electron affinity 18. What would happen if you used an aluminum spoon to stir a solution of FeSO4(aq)? Write a reaction and calculate Eo. 19. Draw an electrochemical cell using Cu and Ag electrodes. 20. 250.0 mL of 0.500 M MnO4- are required to titrate a 100.0 ml sample of SO3-2. Calculate the [SO3-2] 2MnO4- + 3SO32- + H2O → 2MnO2 + 3SO42- + 2OHHow is the breathalyzer reaction used to determine blood alcohol content (you might need to look this up in your textbook? 21. 22. 2H+ + Mg → Mg2+ +H2 Oxidizing agent__________ Worksheet # 9 Reducing agent_________ 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: Anode reaction: Overall reaction: 2. Cathode: Cathode reaction: Voltage: ZnCl2(l) electrolytic cell (electrowinning) Anode: Anode reaction: Cathode: Cathode reaction: Overall reaction: 3. CuSO4(aq) electrolytic cell (electrowinning) Anode: Anode reaction: Overall reaction: 4. Cathode: Cathode reaction: MTV: The electrolysis of 1M NaI (electrowinning) Anode: Anode reaction: Overall reaction: 5. MTV: Cathode: Cathode reaction: MTV: The reaction needed to make Al. The electrolyte is (molten or aqueous). and its phase is 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: Anode reaction: Cathode: Cathode reaction: Possible Electrolyte: 7. The reaction needed to nickel plate a copper penny. Anode: Anode reaction: Cathode: Cathode reaction: Possible Electrolyte: 8. The reaction used in the electrorefining of lead. Anode: Anode reaction: Possible Electrolyte: Cathode: Cathode reaction: 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: Anode reaction: Overall reaction: Cathode: Cathode reaction: Voltage: 2. The electrolytic cell used to produce Al. Phase (aqueous or molten) Cathode: Cathode reaction: 3. Electrolyte: Anode: Anode reaction: Overall reaction: The electrolysis KI(aq) Anode: Anode reaction: Overall reaction: Cathode: Cathode reaction: MTV 4. The electrorefining of Pb Anode: Anode reaction: 5. Cathode: Cathode reaction: Nickel plating a iron nail. Anode: Cathode: Anode reaction: Cathode reaction: Electrolyte The -ve side of the power supply is connected to the 6. Draw an Ag/ Zn electrochemical cell. 7. Draw a KF(l) electrolytic cell. 8. Draw a KF(aq) electrolytic cell. 9. Draw a FeI2(aq) electrolytic cell. 10. Draw a Cd/Pb electrochemical cell. Cd is not on the reduction chart, however, the Cd electrode gains mass and the total cell potential is 0.5 v. Determine the half-cell potential for Cd. 11. Write the overall reaction and describe the anode and cathode for a Zn/C, fuel, alkaline and lead/acid cell. 12. 2HIO3 + 5H2SO3 → oxidizing agent substance reduced I2 + 5H2SO4 + H2O substance oxidized reducing agent 13. What is the electrolyte in a fuel cell? 14. What is the fuel in a fuel cell? 15. Describe the differences and similarities between an electrolytic and electrochemical cell. 16. Describe and give two examples of electrowinning. 17. Describe and give one example of electrorefining. 18. List three metals that can be won from aqueous solution. 19. List three metals that cannot be won from aqueous solution. 20. List the electrolyte in each of the following. Fuel cell, Alkaline battery Dry Cell (Leclanche) Lead acid battery 21. State two metals that can be used to cathodically protect Fe. Describe how they protect iron from corrosion. 22. Write the half reaction that describes the corrosion of iron. 23. Write the half reaction that describes the reduction reaction that occurs when iron corrodes in air and water. 24. Why does iron corrode faster in salt water? 25. Write the anode and cathode reaction in an electrolytic cell with a CaCl2(l) electrolyte. 26. Explain why you would choose Zn or Cu to cathodically protect iron? 27. Choose a suitable redox reactant to oxidize Cl- to ClO4- in a redox titration. 28. Describe as an electrochemical or electrolytic cell: a) Fuel cell b) Charging a car battery c) Discharging a car battery d) Ni plating e) Industrial Al production f) Cl2 production 29. Write the anode and cathode reactions for each of the above processes. 30. Al and AgNO3(aq) are mixed and the surface of the Al darkens. List the two oxidizing agents in decreasing strength. List the two reducing agents in decreasing strength. 31. Analyze This Label each anode and cathode. Write each anode and cathode reaction. Indicate the ion migration in each cell. Determine the initial cell voltage of the electrochemical cell. Determine the MTV for the electrolytic cell. Will electrolysis occur? Indicate electron flow. Indicate all electrodes that gain mass. Indicate all electrodes that lose mass. What happens to [NO3-] in the Mg half-cell? What happens to the [Ag+] in the Ag half-cell? What happens to [Mg2+] in the Mg half-cell? What is the equilibrium electrochemical cell potential? What chemical is made at the Pt electrode on the right? What chemicals are made at the Pt electrode on the left? 1.0 M KNO3 Mg Ag 1 M Mg(NO3)2 Pt 1 M AgNO3 Pt 1 M CuSO4 Quiz #1 Agents, Spontaneous Reactions, Oxidation #’s, and Strength 1. In a redox reaction, the species that loses electrons A. is oxidized B. C. D. 2. Which of the following is the strongest oxidizing agent? A. B. C. D. 3. Fe3+ to Fe2+ Fe2+ to MnO4MnO4- to Fe2+ MnO4- to Mn2+ As an element is oxidized, its oxidation number A. B. C. D. 7. Fe2+ Fe3+ Mn2+ MnO4- MnO4- + 5Fe2+ + 8H+ → Mn2+ + 5Fe3+ + 4H2O During the reaction, electrons transfer from A. B. C. D. 6. Ag, Pt, Au Pt, Au, Ag Au, Ag, Pt Au, Pt, Ag MnO4- + 5Fe2+ + 8H+ → Mn2+ + 5Fe3+ + 4H2O The oxidizing agent in the reaction is A. B. C. D. 5. Cu2+ Pb2+ Ni2+ Sn2+ 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. B. C. D. 4. is called the cathode gains mass at the electrode decreases in oxidation number increases as electrons are lost decreases as electrons are lost increases as electrons are gained decreases as electrons are gained A solution of 1.0 M Pb(NO3)2 will not react with a container made of A. B. C. D. Cu Fe Sn Zn 8. A spontaneous redox reaction occurs when a piece of iron is placed in 1.0 M CuSO4. The reducing agent is A. B. C. D. 9. A substance is oxidized when it A. B. C. D. 10. Hg2+ weaker reducing agent than Cu2+ weaker oxidizing agent than Cu2+ stronger reducing agent than Cu2+ stronger oxidizing agent than Cu2+ The species which gains electrons in a redox reaction A. B. C. D. 13. Ti2+ is a weaker reducing agent than Sn2+ Ti2+ is a weaker oxidizing agent than Sn2+ Ti2+ is a stronger reducing agent than Sn2+ Ti2+ is a stronger oxidizing agent than Sn2+ Consider the following redox reaction : Hg2+ + Cu → Hg + Cu2+ . In this reaction, A. B. C. D. 12. loses protons gains protons loses electrons gains electrons 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. B. C. D. 11. is a Fe Cu2+ H2O SO42- loses mass is oxidized is the oxidizing agent increases in oxidization number 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 1 2 3 4 Ion U3+ V2+ V2+ Y3+ Metal Y U Y V Observation reaction reaction reaction no reaction The oxidizing agents from the strongest to the weakest are A. B. C. D. V2+, U3+, Y3+ U3+, V2+, Y3+ Y3+, U3+, V2+ 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. B. C. D. 2. When NO2 reacts to form N2O4 the oxidation number of nitrogen A. B. C. D. 3. Cu2+ and Fe2+ Pb2+ and Sn2+ Co2+ and Cr2+ Mn2+ and Cr2+ increases by 2 increases by 4 increases by 8 does not change Consider the following redox equation: 12H+(aq) + 2IO3-(aq) + 10Fe2+(aq) → 10Fe3+(aq) + I2(s) + 6H2O(l) The reducing agent is A. B. C. D. 4. The oxidation number of nitrogen increases in A. B. C. D. 5. I2 H+ Fe2+ IO3- NO3- → NO N2O4 → NI3 NH3 → NH4+ NO2 → N2O5 Which of the following represents a balanced reduction half-reaction? A. B. C. D. VO2 + 2H+ + 2e- → V2+ + H2O VO2 + H2 → V2+ + H2O + leVO2 + 2H+ + le- → V2+ + H2O VO2 + 4H+ + 2e- → V2+ + 2H2O 6. Consider the following half reaction: Sb2O3 + 6H+ + 6e- ⇄ 2Sb + 3H2O The oxidation number of antimony in Sb2O3 A. B. C. D. 7. Consider the following unbalanced half-reaction The balanced half-reaction would have A. B. C. D. 8. 10. HClO2 ⇄ HClO 1 electron on the left 1 electron on the right 2 electrons on the left 2 electrons on the right The oxidation number of platinum in Pt(H2O)42+ is A. B. C. D. 9. increases by 3 increases by 6 decreases by 3 decreases by 6 +2 0 +4 +1/2 Consider the following half-reaction: BrO- → BrThe balanced equation for the half-reaction is A. BrO- + 2H+ + 2e- → Br- + H2O B. C. D. BrO- + 2H+ → Br- + H2O + 2eBrO- + H2O → Br- + 2OH- + 2eBrO- + H2O + 2e- → Br- + 2OH- Consider the following redox reaction: 2MnO4- + 5CH3CHO + 6H+ → 5CH3COOH + 2Mn2+ + 3H2O The species that loses the electron is A. B. C. D. 11. H2O MnO4CH3CHO CH3COOH Hydrogen has an oxidation number of –1 in A. B. C. D. H2 NaH H2O KOH (basic) 12. Consider the following: 2NO3- + 4H+ + 2e- → N2O4 + 2H2O This equation represents A. B. C. D. 13. Which of the following half-reactions is balanced? A. B. C. D. 14. IO3- + 6H+ +5e- → I2 + 3H2O IO3- + 6H+ + 4e- →1/2 I2 + 3H2O IO3- + 6H+ → ½ I2 + 3H2O + 5eIO3- + 6H+ + 5e- → ½ I2 + 3H2O Consider the following redox reaction: Al + MnO4- + 2H2O →Al(OH)4- + MnO2 The chemical species being oxidized is A. B. C. D. 15. reduction oxidation neutralization decomposition Al MnO4Al(OH)4MnO2 Consider the following redox reaction: 6H+ + 6I- + ClO3- → 3I2 + 3H2O + ClThe reducing agent is A. B. C. D. 16. Nitrogen has an oxidization number of zero in A. B. C. D. 17. II2 H+ ClO3- N2 NO2 NH3 HNO3 When MnO4- reacts to form Mn2+, the manganese in MnO4- is A. B. reduced as its oxidation number increases reduced as its oxidation number decreases C. D. 18. oxidized as its oxidation number increases oxidized as its oxidation number decreases Consider the following reaction: 2HNO3 + 3H2S → 2NO + 3S + 4H2O The nitrogen in HNO3 undergoes A. B. C. D. 19. reduction oxidation electrolysis neutralization The oxidation number in carbon in CaC2O4 is A. B. C. D. +2 +3 +4 +6 20. Consider the following redox reaction: 2Cr3+(aq) + 3Cl2(aq) + 7H2O(l) → Cr2O72-(aq) + 6Cl-(aq) + 14H+(aq) The species which loses electrons is A. B. C. D. Cl2 Cr3+ H2O Cr2O72- Quiz #3 Balancing Redox reactions- Acid & Base Cell Potentials 1. Consider the following overall reaction: 2Rh+ + Pb(s) → 2Rh(s) + Pb2+The E0 for the half-reaction Rh+ + e- ⇄ A. B. C. D. 2. Rh is -0.86 V -0.60 V +0.60 V +0.86 V 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. B. C. 3. E0 = 0.73 V The reduction potentials of all the half-cells would remain unchanged The reduction potentials of all the half-cells would increase by 0.76 V The reduction potentials of all the half-cells would have positive values Three beakers contain 1.0 M CuCl2. A piece of metal is placed in each of the beakers BEAKER 1 2 3 SOLUTION CuCl2 CuCl2 CuCl2 METAL Zn Ag Ni Reactions occur in A. B. C. D. 4. beaker 2 only beakers 1, 2, and 3 beakers 1 and 2 only beakers 1 and 3 only Consider the following redox reaction: 3SO2 + 3H2O + ClO3- → 3SO42- + 6H+ + Cl- D. The reduct The reduction half-reaction is A. B. C. D. 5. What two substances are produced when Cr and 1.0 M MnO4- react in a basic solution? A. B. C. D. 6. II2 ClCl2 The substances H2O2, H3PO4 and H2SO3 in order of increasing strengths as oxidizing agents are. A B. C. D. 8. Mn2+ and Cr3+ MnO2 and Cr3+ Mn2+ and Cr2+ MnO2 and CrO42- Bromine, Br2, will react spontaneously with A. B. C. D. 7. ClO3- + 6H+ → Cl- + 3H2O + 6eClO3- + 6H+ + 6e- → Cl - + 3H2O SO2 + 2H2O → SO42- + 4H+ + 2eSO2 + 2H2O + 2e- → SO42- + 4H+ H2O2, H3PO4, H2SO3 H2SO3, H3PO4, H2O2 H3PO4, H2SO3 , H2O2 H2O2,H2SO3 , H3PO4 Consider the following overall equation for an electrochemical cell: 3Ag+ + Cr → Cr3+ + 3Ag At standard conditions ,the initial cell voltage is A. B. C. D. +0.06 V +0.39 V +1.21 V +1.54 V 9. A solution of 1.0 M Co(NO3)2 should be stored in a container made of A. B. C. D. 10. A strong oxidizing agent has a A. B. C. D. 11. Zn ClSn2+ Fe3+ Which equation represents a redox reaction? A. B. C. D. 15. spontaneous and Eo is positive. spontaneous and Eo is negative. non-spontaneous and Eo is positive. non-spontaneous and Eo is negative Referring to the data booklet, which of the following can act as an oxidizing agent but not as a reducing agent? A. B. C. D. 14. IO3- and I2 SO42- and S BrO3- and Br AuCl4- and Au Consider the following redox reaction: Co2+(aq) + 2Ag(s) → 2Ag+(aq) + Co(s) The reaction is A. B. C. D. 13. weak attraction for electrons strong attraction for electrons weak ability to become reduced strong ability to become oxidized The two species which react spontaneously in acidic solutions are A. B. C. D. 12. tin zinc aluminum magnesium Pb2+ + 2Cl- → PbCl2 CaO + CO2 → CaCO3 Mg + 2HCl → MgCl2 + H2 HCl + NaOH → NaCl + H2O In a redox reaction, ClO- was converted to Cl- in a basic solution. The balanced half-reaction for this process is A. B. ClO- + H2O + 2e- → Cl- + 2OHClO- + 2OH- → Cl- + 2e- + H2O C. D. ClO- + H2O → Cl- + 2e- + 2OHClO- + 2OH- + 2e- → Cl- + H2O Quiz #4 Electrochemical Cells/Electrolytic Cells voltmeter 1.0 M KNO3 Zn Pb 1M Zn(NO3)2 1. In the electrochemical call above, the electrons flow from A. B. C. D. 2. zinc to lead and the mass of zinc increases zinc to lead and the mass of lead increases lead to zinc and the mass of zinc increases lead to zinc and the mass of lead increases The initial cell voltage is A. B. C. D. 3. 1M Pb(NO3)2 -0.89 V -0.63 V +0.63 V +0.89 V In an operating lead-zinc electrochemical cell shown above, the cathode A. B. C. D. gains mass as anions are reduced loses mass as anions are reduced gains mass as cations are reduced loses mass as cations are reduced 4. The equation for the half-reaction at the anode is A. B. C. D. 5. Zn2+ + 2e- → Zn Pb2+ + 2e- → Pb Zn → Zn2+ + 2ePb → Pb2+ + 2e- The equation for the half-reaction at the cathode is A. B. C. D. 6. Zn2+ + 2e- → Zn Pb2+ + 2e- → Pb Zn → Zn2+ + 2ePb → Pb2+ + 2e- The direction of electron flow in an electrochemical cell is from A. B. C. D. 7. anode to cathode through the external wire cathode to anode through the external wire anode to cathode through the external wire and back through the salt bridge cathode to anode through the external wire and back through the salt bridge Which of the following is formed at the anode during the electrolysis of 1.0 M NaI? A. B. C. D. 8. I2 O2 H2 Na As this cell operates A. B. C. D. Cl- is oxidized at the anode Mg2+ is oxidized at the anode Cl- is oxidized at the cathode Mg2+ is oxidized at the cathode Power Source Pt 9. + Pt In an operating electrochemical cell, the anions migrate A. B. Molten 2 through the wire towards theMgCl anode towards the cathode through the wire C. D. towards the anode through the salt bridge towards the cathode through the salt bridge voltmeter 1.0 M KNO3 Mg Cu 1 M Mg(NO3)2 10. As the above electrochemical cell operates A. B. C. D. 11. 14. Cu → Cu2+ + 2eCu2+ + 2e- → Cu Mg → Mg2+ + 2eMg2+ + 2e- → Mg In the above electrochemical cell, the initial voltage is A. B. C. D. 13. nitrate ions migrate into the copper half-cell copper(II) ions migrate through the salt bridge magnesium ions migrate through the salt bridge potassium ions migrate into the magnesium half-cell In the above electrochemical cell, the reaction at the anode is A. B. C. D. 12. 1 M CuSO4 2.03 V 2.52 V 2.71 V 2.89 V Which of the following aqueous solutions produces H2(g) and O2(g) during electrolysis A. 1.0 M KI B. 1.0 M CuI2 C. 1.0 M K2SO4 D. 1.0 M CuSO4 In the electrolysis of molten zinc chloride, the half-reaction at the anode is A. B. C. D. Cl2 + 2e- → 2Cl2Cl- → Cl2 + 2eZn2+ 2e- → Zn Zn → Zn2+ + 2e- voltmeter 1.0 M KNO3 Ni Ag 1 M Ni(NO3)2 15. The initial cell voltage at 25oC is A. B. C. D. 16. -1.06 V -0.54 V +0.54 V +1.06 V The balanced equation for the overall reaction is A. B. C. D. 17. 1 M AgNO3 Ni+(aq) + Ag(s) → Ag+(aq) + Ni(s) Ni(s) + Ag+(aq) → Ag(s) + Ni+(aq) Ni2+(aq) + 2Ag(s) → 2Ag+(aq) + Ni(s) Ni(s) + 2Ag+(aq) → 2Ag(s) + Ni2+(aq) This redox reaction occurs because A. B. C. D. Ag(s) is a stronger oxidizing agent than Ni(s) Ag(s) is a weaker reducing agent than Ni(s) Ag+(aq) is a stronger reducing agent than Ni2+(aq) Ag+(aq) is a weaker oxidizing agent than Ni2+(aq) voltmeter 1.0 M KNO3 Au Pb 18. The direction of the electron flow is A. B. C. D. 19. As the cell operates A. B. C. D. 20. -1.37 V 0.00 V 1.37 V 1.63 V Which of the following is a balanced half-reaction in base? A. B. C. D. 22. NO3- and K+ will migrate toward the Pb half-cell NO3- and K+ will migrate toward the Au half-cell NO3- migrates toward the Pb half-cell and K+ will migrate toward the Au NO3- migrates toward the Au half-cell and K+ will migrate toward the Pb The initial voltage is A. B. C. D. 21. from Au to Pb through the wire from Pb to Au from the wire from Au to Pb through the salt bridge from Pb to Au through the salt bridge Cl2 + 3H2O → ClO3- + 6H+ + 5eCl2 + 6OH- → ClO3- + 5e- + 3H2O Cl2 + 6H2O → 2ClO3- + 12H+ + 10eCl2 + 12OH- → 2ClO3- + 6H2O + 10e- In which of the following unbalanced equations does chromium undergo oxidation? A. B. C. D. Cr3+ → Cr Cr3+ → Cr2+ Cr3+ → Cr2O72CrO42- → Cr2O72- Quiz #5 Application of Cells 1. The corrosion of iron can be prevented by attaching a piece of zinc to the iron because A. B. iron acts as an anode zinc reduces more readily than iron the C. D. 2. An iron spoon is electroplated with copper. The equation representing the reduction reaction is A. B. C. D. A. B. C. D. 4. ANODE carbon pure lead pure lead impure lead 3. In an operating zinc-copper electrochemical cell, the oxidizing agent A. B. C. D. loses electrons at the anode loses electrons to the cations gains electrons at the cathode gains electrons from the anions extraction of aluminum from bauxite purification of lead from an impure anode recovery of zinc from a zinc sulphide solution production of chlorine from a sodium chloride solution oxidation of anions reduction of cations reduction at the anode oxidation at the cathode Hydrogen and oxygen react to provide energy in a A. B. C. D. 7. CATHODE impure lead carbon impure lead pure lead Electroplating always involves the A. B. C. D. 6. Cu2+(aq) + 2e- → Cu(s) Cu(s) → Cu2+(aq) + 2eFe2+(aq) + 2e- → Fe(s) Fe(s) → Fe2+(aq) + 2e- An example of electro refining is the A. B. C. D. 5. electrons flow from the zinc to the iron iron ions form more readily than zinc ions dry cell fuel cell alkaline cell lead-acid storage cell En electrolytic process is used to purify impure lead. The electrodes are 8. In the cell below the half-reaction at the cathode is Cu2+ + 2e- → Cu(s) 2SO42- → S2O82- + 2eH2O → ½ O2(g) + 2H+ + 2e2H2O + 2e- → H2(g) + 2OH- A. B. C. D. Power Source - + Iron Key Pt 1.0 M CuSO4 9. In the electrolysis of molten PbBr2, the products at the anode and cathode are CATHODE (INERT) A. B. C. D. 10. ANODE (INERT) Br2 O2 Pb Br2 H2 Pb Br2 Pb Under which conditions could an electrochemical cell provide 0.93V? Cathode A. B. C. D. Anode Cu Mg Mg Cu Ag Pb 11. The reduction reaction in the above electrochemical cell is A. B. C. D. 12. Ba2+ Al3+ Sn2+ Na+ produce fuel electrolyze fuel produce hydrogen produce electricity If a piece of nickel is to be gold-plated using an electrolytic process, which half-reaction occurs at the cathode? A. B. C. D. 17. the nickel coin must be the cathode the cathode must be made of copper the electrons must flow to the anode the solution must contain nickel ions The principal function of a fuel cell is to A. B. C. D. 16. water forming oxygen gas water forming hydrogen gas sea water forming chlorine gas sea water forming bromine liquid Which of the following ions can be reduced from an aqueous solution A. B. C. D. 15. 2e- → Pb Pb2+ + 2ee- → Ag Ag+ + e- To plate a nickel coin with copper A. B. C. D. 14. Pb2+ + Pb → Ag+ + Ag → An industrial process involving electrolysis is the reduction of A. B. C. D. 13. Pb Ag Ni → Ni2+ + 2eNi2+ + 2e- → Ni Au → Au3+ + 3eAu3+ + 3e- → Au Consider the following redox reaction As2O3 + 2NO3- + 2H2O + 2H+ → 2H3AsO4 + N2O3 In this reaction, nitrogen A. B. C. D. 18. In an electrochemical cell, the cathode A. B. C. D. 19. Quiz #6 Which of the following metals could be used to cathodically protect a sample of lead? iron gold silver copper A piece of iron can be prevented from corroding by A. B. C. D. 3. 2I- → I2 + 2eNa+ + e- → Na H2O + ½ O2 + 2H+ + 2e2H2O +2e- → H2 + 2OH- Corrosion & Cathodic Protection Titration A. B. C. D. 2. is reduced loses mass is the reducing agent is the site of reduction 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. B. C. D. 1. loses electrons and increases in oxidation number gains electrons and increases in oxidation number loses electrons and decreases in oxidation number gains electrons and decreases in oxidation number making it a cathode placing it in an acidic solution attaching a small piece of lead to it attaching a small piece of gold to it To determine the [Fe2+] in a solution of FeSO4 by a redox titration, a suitable reagent would be an acidified solution of A. B. C. Cr3+ Mn2+ SO42- D. 4. As a metal corrodes, A. B. C. D. 5. less than 0.34 V greater than 1.50 V greater than 0.85 V but less than 1.50 V greater than 0.34 V but less than 0.85 V Consider the following redox equation: Br2 + SO2 + Na2SO4 + 2H2O → 2H2SO4 + 2NaBr Which of the following is gaining electrons? A. B. C. D. 9. Mn Cu Pb Sn A student attempted to determine the Eo (volts) of the following half-reaction: Pd2+ + 2e- → Pd Pd2+ reacts with Cu(s) but not with Hg(l). Based on the above, the Eo (volts) of a Pd half-cell is A. B. C. D. 8. Paint the iron Cover the iron with grease Attach a piece of lead tot he iron Attach a piece of magnesium to the iron Corrosion of iron can be prevented by attaching a piece of A. B. C. D. 7. it gains electrons it becomes reduced it acts as a reducing agent its oxidation number decreases Which method will cathodically protect a piece of iron? A. B. C. D. 6. Cr2O72- Br2 SO2 H2O Na2SO4 The reaction that occurs when pieces of lead, zinc, copper and silver are placed in a solution of Ni(NO3)2 is A. B. C. D. Pb + Ni2+ → Pb2+ + Ni Zn + Ni2+ → Zn2+ + Ni Cu + Ni2+ → Cu2+ + Ni 2Ag + Ni2+ → 2Ag+ + Ni voltmeter 1.0 M KNO3 Pb Cu 1M Pb(NO3)2 10. In the electrochemical cell above, the electrons flow from A. B. C. D. 11. copper to lead through the wire lead to copper through the wire copper to lead through the salt bridge lead to copper through the salt bridge In the electrochemical cell above, the initial Eo value is A. B. C. D. 12. 1M Cu(NO3)2 0.03 V 0.21 V 0.29 V 0.47 V A reaction that occurs during the corrosion of iron is Fe + 3e- → Fe3+ Fe → Fe2+ + 2eFe2+ + 2e- → Fe Fe3+ + e- → Fe2+ A. B. C. D. 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. B. C. D. 0.0131 L 0.0262 L 0.200 L 0.400 L Redox Web Review 1) Which most readily gains electrons? Cu 2) Cu2+ Fe2+2 Zn2+ Au3+ Which most readily loses electrons? Hg(l) Cu2+ Sn4+ Ba Al Calculate the cell potentials or voltages (E0) Indicate spontaneity. 3. Cl2 + 2Br- → 2Cl- +Br2 4. 2MnO4- + 5Pb +16H+ 5. Will AgNO3 react with Zn? Write a balanced redox reaction and calculate Eo 6. → 2Mn2+ + 8H2O + 5Pb2+ What would happen if you used an iron spoon to stir a solution of Al2(SO4)3(aq) ? Write a balanced redox reaction and calculate Eo. 7. What are the differences between an electrochemical cell and an electrolytic cell? Electrochemical cell 8. Electrolytic cell 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 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 Electrolytic Cell 10. Draw an operating electrochemical cell using an Al half-cell and a Mg half-cell. Label the parts of the electrochemical cell including the anode or cathode, and all reagents and materials used. Write the reactions and determine the E0. 11. Write the half reaction that occurs at each electrode during the electrolysis of aqueous 1.0 M NaI. Anode : Cathode : What is the minimum required voltage for this process? 12. Write the half reaction that occurs at each electrode during the electrolysis of molten NaI. Anode : Cathode : What is the minimum required voltage for this process? 13. Aluminum is produced industrially from aluminum oxide, Al2O3. Demonstrate your understanding of this process by (i) Describing how the process is carried out, (ii) Writing equations of the reactions involved in the process, and (iii) Describing how the problem of the high melting point ofAl2O3 is overcome. 14. Consider the following redox data: 3V + 2Ga3+ → 3V2+ + 2Ga Eo = +0.64 V 3V2+ + 2Al Eo = +0.46 V → 3V + 2Al3+ 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) 16. HV2O73- Balance the following redox reaction occurring in basic solution: MnO4- 17. → + C2O42- → MnO2 + CO2 250.0 ml 0.200M MnO4- reacts with excess SO32-. How many grams of MnO2 are produced? 2MnO4- + 3SO32- +H2O → 2MnO2 +3SO42- + 2OH- 18. Determine the oxidation number for each bold atom. MnO2 IO3- Cr2O72- C2O42- HOOH NO3- H3PO4 Na2C2O4 19. Al(NO3)3 NH4Cl I2 N2O3 NaH Pt(H2O)42+ 250.0 mL of 0.500M MnO4- are required to titrate a 100.0 ml sample of SO3-2. Calculate the [SO3-2] 2MnO4- + 3SO32- +H2O → 2MnO2 + 3SO42-+ 2OH- 20. How is the breathalyzer reaction used to determine BAC? Write the reaction and describe how it works. 21. 2H+ + → Mg Mg2+ + H2 Determine the Oxidizing agent__________ and the Reducing agent_________ 22. Choose a suitable redox reactant to oxidize Cl- toClO4- in a redox titration. 23. Describe as an electrochemical or electrolytic cell: a) Fuel cell b) Charging a car battery c) Discharging a car battery d) Ni plating e) Industrial Al production f) Cl2 production g) Electrowinning 24) Which of the reactants is gaining electrons? Which of the reactants is the oxidizing agent? Br2 + SO2 + Na2SO4 + H2O → 2H2SO4 + 2NaBr 25) A student studied the following reactions and she recorded: Pd2+ + Cu → Pd + Cu2+ spontaneous Pd2+ + Au → no reaction Pd2+ + Hg → no reaction Au3+ + Hg → Au + Hg2+ spontaneous List the oxidizing agents from strongest to weakest. List the reducing agents from strongest to weakest. Predict if the reaction will occur. Au3+ 26) + Cu → 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 oxidation of water oxidation of water in acid reduction of water reduction of water in alkaline oxidation of H2 in water oxidation of H2 in acid oxidation of H2 in base reduction of Cr2O72- in acid reduction of HBr Eo 35. Completely analyze the following electrochemical cell. voltmeter 1.0 M KNO3 Zn Cu 1 M Zn(NO3)2 1 M Cu(NO3)2 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. voltmeter 1.0 M KNO3 H2(g) Cu 1 M HCl 1 M Cu(NO3)2 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. Power Source - + C C Molten Al2O3 Anode Reaction Cathode Reaction Chemicals produced at the anode Chemicals produced at the cathode The electrons flow from __to __ The chemical used to lower the mp is: Which electrode is the anode ? 38. Completely analyze the following electrolytic cell. Note that the electrodes are not inert and because of that, the anode might oxidize. Power Source - + Cu Cu 1 M NaF 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. B. C. D. Fe CO CO2 Fe2O3 2. Consider the following reaction: 2HNO2 + 2I- + 2H+ → 2NO + I2 +2H2O The oxidation number for each nitrogen atom A. B. C. D. 3. Which of the following reactions is spontaneous? A. B. C. D. 4. increases by 1 increases by 2 decreases by 1 decreases by 2 2I- + Ag → Ag+ + I2 Co2+ + Cu → Co + Cu2+ Cu2+ + Pb → Pb2+ + Cu Ni2+ + 2Ag → 2Ag+ + Ni Consider the following redox reaction for a lead-acid storage cell: Pb + PbO2 + 4H+ + 2SO42- → 2PbSO4 + 2H2O The balanced, reduction half reaction is A. B. C. D. 5. -1.12 V -0.40 V +0.40 V +1.12 V Which of the following involves a nonspontaneous redox reaction? A. B. C. D. 7. SO42- → 2PbSO4 + 2e2H+ + SO42- → PbSO4 + 2H2O + 2e+ 4H+ + SO42- + 2e- → PbSO4 + 2H2O + 2SO42 + 2H2O + 2e- → PbSO4 + 2OH- 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. B. C. D. 6. Pb + Pb + PbO2 PbO2 fuel cell electroplating redox titration carbon dry cell 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. C. D. 8. Which of the following is not a redox reaction? A. B. C. D. 9. Cathode O2 Na Cl2 Cl2 H2 Cl2 H2 Na 1.0 M H+ 1.0 M Ag+ 1.0 M Sr2+ 1.0 M Mn2+ Which of the following half-reactions are balanced? A. B. C. D. 13. Anode A solution containing an unknown cation reacts spontaneously with both zinc and copper. The unknown cation is A. B. C. D. 12. 0.26 V 0.28 V 0.54 V 0.80 V What substances are formed at the anode and cathode during electrolysis of molten sodium chloride? A. B. C. D. 11. Cu + Br2 → CuBr2 CO + H2O → CO2 + H2 CH4 + H2O → CO2 + 2H2O NaOH + HCl → NaCl + H2O What is the minimum voltage required to form nickel from an aqueous solution of NiI2 using inert electrodes? A. B. C. D. 10. 0.060 moles 1.5 moles 0.30 moles ClO- + H2O + e- → Cl2 + 2OH2ClO- + H2O + 2e- → Cl2 + 3OH2ClO- + 2H2O + 2e- → Cl2 + 4OH2ClO- + 2H2O → Cl2 + 4OH- + 2e- Which of the following is a spontaneous redox reaction? A. B. C. D. 14. Ag+ + I- → AgI Ag+ + Fe2+ → Ag + Fe3+ 3Ag+ + Au → 3Ag + Au3+ 2Ag+ + Ni2+ → 2Ag + Ni Salting the roads during the winter increases the amount of corrosion of cars. The is because the salt A. B. C. D. reacts with the iron provides an electrolyte acts as a reducing agent acts as an oxidizing agent Consider the following electrochemical cell for the next five questions. voltmeter 1.0 M KCl Cuu Ni 1M Cu(NO3)2 15. The half-reaction that occurs at the anode is A. B. C. D. 16. 0.41 V 0.78 V 0.34 V 0.60 V The following ions migrate to the Cu electrode A. B. C. D. 19. Ni → N2+ + 2eNi2+ + 2e- → Ni Cu → Cu2+ + 2eCu2+ + 2e- → Cu The cell potential or Eo is A. B. C. D. 18. Ni → N2+ + 2eNi2+ + 2e- → Ni Cu → Cu2+ + 2eCu2+ + 2e- → Cu The half-reaction that occurs at the cathode is A. B. C. D. 17. 1M Ni(NO3)2 K+ Cu2+ ClCl- Cu2+ Ni2+ Ni2+ NO3NO3- 2e- The electrons flow A. B. through the salt bridge from Cu to Ni through the salt bridge from Cu to Ni C. D. 20. Which of the following will not react spontaneously with 1.0 M HCl? A. B. C. D. 21. Zn → Zn2+ + 2eZn2+ + 2e- → Zn 2Cl- → Cl2 + 2eCl2 + 2e- → 2Cl- In the electrolysis of molten zinc chloride, the half-reaction at the anode is A. B. C. D. 25. a voltmeter. a salt bridge. a power supply. an aqueous solution. In the electrolysis of molten ZnCl2 using carbon electrodes, the reaction that occurs at the anode is A. B. C. D. 24. nickel sodium aluminum magnesium In order for an electrolytic cell to operate, it must have A. B. C. D. 23. tin lithium mercury magnesium Which of the following can be produced by electrolysis from a 1.0 M aqueous solution containing its ion? A. B. C. D. 22. through the wire from Cu to Ni through the wire from Ni to Cu Cl2 + 2e- → 2Cl2Cl- → Cl2 + 2eZn2+ + 2e- → Zn Zn → Zn2+ + 2e- The corrosion of iron can be prevented by attaching a piece of A. B. C. D. Mn Cu Pb Sn 26. The oxidation number of carbon in CaC2O4 is A. B. C. D. 27. +2 +3 +4 +6 To plate a nickel coin with copper, A. B. C. D. the nickel coin must be the cathode. the cathode must be made out of copper the electrons must flow to the anode the solution must contain nickel ions Consider the following electrochemical cell for the next five questions. voltmeter 1.0 M KNO3 Zn Cuuu 1M Cu(NO3)2 28. Which of the following statements apply to this electrochemical cell? I II III 29. Electrons flow through the wire toward the copper electrode. The copper electrode increases in mass. 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 The balanced equation for the overall reaction is A. B. C. D. 30. 1M Zn(NO3)2 Zn + Cu2+ → Cu + Zn2+ Cu + Zn2+ → Zn + Cu2+ Zn2+ + Cu → Cu2+ + Zn Cu + Zn → Zn + Cu At equilibrium the voltage of the above cell is A. B. C. -1.10 V 0.00 V +0.42 V D. 31. This redox reaction occurs because A. B. C. D. 32. KCl NaOH H2SO4 KOH The electrolyte used in an automobile battery is A. B. C. D. 37. reduced as its oxidation number increases reduced as its oxidation number decreases oxidized as its oxidation number increases oxidized as its oxidation number decreases The electrolyte used in the alkaline battery is A. B. C. D. 36. spontaneous and Eo is positive spontaneous and Eo is negative non-spontaneous and Eo is positive non-spontaneous and Eo is negative When MnO4- reacts to form Mn2+, the manganese in MnO4- is A. B. C. D. 35. -1.10 V +1.10 V +0.91 V +0.86 V Consider the following redox reaction: Co2+(aq) + 2Ag(s) ⇋ 2Ag+(aq) + Co(s) The reaction is A. B. C. D. 34. Zn is a stronger oxidizing agent than Cu Zn is a stronger reducing agent than Cu, Cu is a stronger oxidizing agent than Zn Zn2+ is a weaker reducing agent than Cu2+ The initial cell voltage at 25 oC is A. B. C. D. 33. +1.10 V KCl NaOH H2SO4 KOH The anode used in the commercial production of Aluminum is A. B. C. D. 38. C Pt Al Al2O3 The anode and cathode used in the electrorefining of impure lead to pure lead are A. B. C. D. 39. Anode Cathode Pure Pb Impure Pb Pb2+ Pb Impure Pb Pure Pb Pb Pb2+ The anode in the LeClanche or common dry cell is A. B. C. D. C Zn Mg KOH 40. Which of the following are electrolytic cells I II III IV Electro winning Electroplating Charging a car battery Fuel cell A. B. C. D. I and II only I, II, and III only II and II only I, II, III, and IV Subjective 1. Balance the following in basic solution. MnO4- 2. 3. + C2O42- → MnO2 + CO2 (basic) 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. Consider the following diagram for the electro refining of lead. Power Source 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 Final burette reading 6. 2.00 13.80 13.80 24.55 24.55 35.32 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 are not inert and can oxidize: Cl-, or Cu will oxidize) 8. a) Identify the product at the anode. b) Write the equation for the reduction half-reaction. c) Write the equation for the overall reaction. Completely analyze the following electrochemical cell. voltmeter 1.0 M KNO3 Snnnn Mn electrodes Chemistry 12 Electrochemistry Practice Test 2 voltmeter 1.0 M KCl Pd Ni 1M Pd(NO3)2 1M Ni(NO3)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 B C D 2. + 2e+ 2e→ Pb → Ni As the cell operates, A B C D 3. Pd → Pd2+ Ni → Ni2+ Pd2+ + 2eNi2+ + 2e- both the K+ and the NO3- migrate into the nickel half-cell both the K+ and the NO3- migrate into the palladium half-cell the K+ migrates into the nickel half-cell and the NO3- migrates into the palladium halfcell the K+ migrates into the palladium half-cell and the NO3- migrates into the nickel halfcell The initial cell voltage is 1.21 V. The reduction potential of Pd2+ is A B C D -1.21 V -.95 V +0.95 +1.21 V 4. What substances are formed at the anode and cathode during electrolysis of molten sodium chloride, NaCl(l)? A B C D 5. Anode Cathode O2 Na Cl2 Cl2 H2 Cl2 H2 Na Consider the following electrolytic cell: Power Source + - Inert Electrode Inert Electrode Molten NaI(l) In the cell above A B C D 6. I- migrates to the anode and gains electrons I- migrates to the cathode and loses electrons Na+ migrates to the anode and loses electrons Na+ migrates to the cathode and gains electrons 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 B C D I and II only I and III only II and III only I, II, and III 7. A fuel cell consumes H2 and O2 gas, uses a KOH electrolyte, and produces electricity. reaction at the anode is A B 2H+ + 2e- → H2 1/2O2 + 2H+ + 2e- → H2O The C D 8. 4OH- → O2 + 2H2O + 4eH2 + 2OH- → 2H2O + 2e- A student investigating redox reactions recorded the following results: V2+ + Te2- → no reaction U4+ + Te2- → U3+ + Te Based on these results, the strengths of the oxidizing agents, arranged from weakest, are A B C D V2+ U4+ U3+ V2+ Te Te Te2Te2- U4+ V2+ V2+ U3+ 9. What is the minimum voltage required to form nickel from an aqueous solution of using inert electrodes? A B C D 0.26 V 0.28 V 0.54 V 0.80 V strongest to NiI2 10. voltmeter 1.0 M KNO3 Zn Ni 1M Zn(NO3)2 1M Ni(NO3)2 Which of the following occurs as the cell operates? A B C D 11. Which of the following reactants would produce an E0 of +0.63 V? A B C D 12. Ag+ + I2 Pb2+ + Zn Mg2+ + Ca Zn2+ + Mn The concentration of Fe2+(aq) can be determined by a redox titration using A B C D 13. the Zn electrode is reduced and increases in mass the Zn electrode is reduced and decreases in mass the Zn electrode is oxidized and increases in mass the Zn electrode is oxidized and decreases in mass KBr SnCl2 KMnO4 (basic) KBrO3 (acidic) Which of the following will oxidize Fe2+? A B C D I2(s) Ni(s) Zn(s) Br2(l) 14. The oxidation number of carbon in C2O42- is A B C D 15. +3 +4 +5 +6 Consider the following reaction: 3As2O3 + 4NO3- + 7H2O → 6H3AsO4 + 4NO The oxidizing agent is A B C D 16. When W2O5 is converted to WO2 in a redox reaction, the W has been A B C D 17. H+ H2O NO3AsO3 reduced since its oxidation number has increased reduced since its oxidation number has decreased oxidized since its oxidation number has increased oxidized since its oxidation number has decreased Consider the following: I II III Water Oxygen gas Nitrogen At 25oC, a piece of iron rusts in the presence of A B C D 18. Which of the following represents a redox reaction? A B C D 19. I only III only I and II only II and III only H2CO3 → H2O + CO2 CuS + H2 → H2S + Cu AgNO3 + NaCl → AgCl + NaNO3 2HCl + Na2SO3 → 2NaCl + H2SO3 The following reaction occurs in an electrochemical cell: 3Cu2+ + Cr → 2Cr3+ + 3Cu The Eo for the cell is A B C D 20. During the corrosion of magnesium, the anode reaction is A B C D 21. Zn → Zn2+ + 2e2Cl- → Cl2 + 2eCl2 + 2e- → 2ClZn2+ + 2e- → Zn Which of the following represents a redox reaction? A B C D 23. Mg → Mg2+ + 2eMg2+ + 2e- → Mg 4OH- → O2 + 2H2O + 4eO2 + 2H2O + 4e- → 4OH- A molten binary salt, ZnCl2, undergoes electrolysis. The cathode reaction is A B C D 22. 0.40 V 0.75 V 1.08 V 2.50 V CaCO3 → CaO + CO2 SiCl4 + 2Mg → Si + 2MgCl2 2NaOH + H2SO4 → 2H2O + Na2SO4 AgBr + 2S2O32- → Ag(S2O3)23- + Br- The process of applying an electric current through a cell to produce a chemical is called A B C D corrosion ionization hydrolysis electrolysis change 24. A student investigating redox reactions recorded the following results: V2+ + Te2- → no reaction U4+ + Te2- → U3+ + Te Based on these results, the strengths of the oxidizing agents, arranged from weakest, are A B C D 25. 27. I2 Cu H2S Ag2S A B C 2Cl- → Cl2 + 2e2H+ → H2 + 2eBrO3- + 6H+ + 5e- → ½ Br2 + 3H2O D BrO3- + 6H+ → ½ Br2 + 3H2O + 5e- Which of the following is not a redox reaction? Cu + Br2 → CuBr2 CO + H2O → CO2 + H2 CH4 + O2 → CO2 + 2H2O NaOH + HCl → NaCl + H2O During the electrolysis of 1.0 M Na2SO4, the reaction at the cathode is A B C D 29. U4+ V2+ V2+ U3+ Consider the redox reaction: 2BrO3- + 10Cl- + 12H+ → Br2 + 5Cl2 + 6H2O the oxidation half-reaction ivolved in this reaction is A B C D 28. Te Te Te2Te2- A spontaneous redox reaction occurs when Sn2+ is mixed with A B C D 26. V2+ U4+ U3+ V2+ Na+ + 1e- → Na 2SO42- → S2O82- + 2e2H2O → O2 + H+ + 4e2H2O + 2e- → H2 + 2OH- An oxidizing agent will cause which of the following changes? A PtO2 → PtO strongest to B C D 30. PtO3 → PtO2 Pt(OH)2 → Pt Pt(OH)22+ → PtO3 Consider the overall reaction of the nickel-cadmium battery: NiO2(s) + Cd(s) + 2H2O(l) → Ni(OH)2(s) + Cd(OH)2(s) Which of the following occurs at the anode as the reaction proceeds? A B C D Cd loses 2e- and forms Cd(OH)2(s) Cd gains 2e- and forms Cd(OH)2(s) NiO2 loses 2e- and forms Ni(OH)2(s) 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 B C D 32. In the electrolysis of molten ZnCl2 using carbon electrodes, the reaction that occurs at the anode is A B C D 33. nickel sodium aluminum magnesium Zn → Zn2+ + 2eZn2+ + 2e- → Zn 2Cl- → Cl2 + 2eCl2 + 2e- → 2Cl- In order for the electrolytic cell to operate, it must have A B C D a voltmeter a salt bridge a power supply 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 Pt2+ reaction b) 2. no reaction reaction List the oxidizing agents in order of strongest to weakest Consider the following reaction for the formation of rust: Fe(s) + ½ O2(g) + H2O(l) → Fe(OH)2 Describe and explain two methods, using different chemical principles, to prevent formation of rust. a) b) 3. Consider the following redox reaction: H2Se + SO42- + 2H+ → Se + H2SO3 + H2O Calculate the Eo for the reaction. 4. Balance the following redox reaction in basic solution: Au + Cl- + O2 → AuCl4- + OH- the 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 7. 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. Consider the following diagram in the electro refining of lead: Power Source Pure Pb Impure Pb a) On the diagram above, label the anode and cathode. b) Write the formula for a suitable electrolyte. c) Write the equation for the reduction half-reaction. d) Write the anode reaction Oxidation and Reduction Reactions Workbook Notes- double click on the lesson number and download Power Point Viewer if you do not have it. Worksheets 1. Oxidation, Reduction, Agents, & Reactions. Quiz WS 1 2. Lab: The Strength of Oxidizing Agents. 3. Oxidation Numbers Spontaneous Reactions WS 2 4. Oxidation Numbers, Application to Reactions. 1 WS 3 5. Balancing Redox Half Reactions Acid/Base. WS 4 6. Balancing Redox Reactions in Acid/Base. WS 5 7. Standard Potentials Using Chart. WS 6 8. Electrochemical Cells. WS 7 2 3 9. Electrochemical Cells Lab. 10. Electrolytic Cells. WS 8 4 12. Application of Electrolytic Cells. WS 9 5 13. Application of Electrochemical Cells: Bat & Cor. WS 10 6 14. Breathalyzer and review. Internet Review Quizmebc 15. Review Practice Test # 1 16. Review Practice Test # 2 11. Electrolytic Cells Lab. 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 2. Reduction 3. Oxidizing agent 4. Reducing agent - loss of electrons - gain of electrons - causes oxidation by undergoing reduction - causes reduction by undergoing oxidation Write half reactions for each of the following atoms or ions. Label each as oxidation or reduction. Al3+ 5. Al -----------> 6. S + 7. 2O2- ----------> 8. Ba2+ 9. 2N3- ----------> 10. Br2 11. P + 12. Ca -----------> 13 Ga3+ + 14. S + 15. H2 ---------> 16. 2H+ + 17. 2F- ----------> F2 18. P3- ----------> P + 2e- ---------> + O2 3e- + S2- reduction 4e- + oxidation 2e- -----------> Ba N2 3e- reduction 6e- oxidation 2Br- reduction + 2e- ---------> + P3- ----------> Ca2+ reduction 2e- + 3e- -----------> Ga 2e- ---------> 2H+ 2e- ---------> 2e- H2 + oxidation reduction S2+ oxidation reduction oxidation reduction 2e3e- oxidation oxidation Balance each spontaneous redox equation. Identify the entities reduced and oxidized. State the reducing agent and the oxidizing agent. 19. Al & Zn2+ 2Al + oxidized reducing agent 20. F2 & & & 3Zn 2O2→ oxidized reducing agent 4F- O2 + Ca 2Ca + oxidized reducing agent 22. Al3+ 2Al3+ + O2- 2F2 + reduced oxidizing agent 21. O2 3Zn2+ → reduced oxidizing agent O2 → reduced oxidizing agent 2Ca2+ + 2O2- + 3Li+ Li Al3+ + reduced oxidizing agent 3Li → oxidized reducing agent Al Label the species that is reduced, that is oxidized, the reducing agent and the oxidizing agent. Fe2+ 23. Co Fe Co Co2+ + 2e- oxidation + Ni → Ni2+ + 2e- oxidation + Pb → Co2+ Fe2+ + Fe + 2e- + 3 Ag + 1e- + Cu → reduction 3 Ag+ 24. Ni Ag 25. → → + → Ni3+ Ag+ reduction Cu2+ Pb2+ → Pb Cu 26. → + 2e- oxidation + 2 Sn → Sn2+ + 2e- oxidation + 2 F- → F2 + 2e- oxidation Cu2+ + 2e- + 2 Sn2+ O2 + + F2 Co2+ + Sn → 2O2- → O2- → Co → Coreduction 28. List the species (formulas from above) that lose electrons: 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- + 4e- reduction Co2+ 2F- Co → reduction O2 27. Pb2+ Cl2 ----------> Substance oxidized Oxidizing agent ICl2 Cl- + I2 Reducing agent Substance reduced ICl2 2e- 31. Co + Fe3+ Substance oxidized Oxidizing agent 32. Cr6+ + Co2+ -----------> Fe2+ Co Fe3+ Fe2+ Cr6+ Fe2+ Reducing agent Substance reduced Co Fe3+ Cr3+ Fe3+ -----------> Substance oxidized Oxidizing agent + + Reducing agent Substance reduced Fe2+ Cr6+ Redox Half Reactions and Reactions WS #2 1. State the Oxidation Number of each of the elements that is underlined. a) NH3 c) ZnSO3 e) Na g) AgNO3 i) SO2 k) Ca(ClO3)2 m) HPO32o) MnO2 -3 4 0 5 4 5 3 4 b) H2SO4 d) Al(OH)3 f) Cl2 h) ClO4j) K2Cr2O4 l) K2Cr2O7 n) HClO p) KClO3 6 3 0 7 3 6 1 5 q) PbO2 s) K2SO4 u) Na2O2 w) Fe2O3 y) NaIO3 aa) NO3cc) CaH2 ee) Fe(H2O)63+ 4 6 -1 3 5 5 -1 +3 r) PbSO4 t) NH4+ v) FeO x) SiO44z) ClO3bb) Cr(OH)4 dd) Pt(H20)5(0H)2+ ff) CH3COOH 2 -3 2 -2 5 4 +3 0 2. What is the oxidation number of carbon in each of the following substances? a) CO c) CO2 e) C2H6 2 b) C 2- 4 -3 0 d) CO3 f) CH3OH 4 -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) Substance oxidized --------> Zn Cu (s) + Zn2+ (aq) Substance reduced Cu2+ Cu2+ Oxidizing agent Reducing agent Zn b) Cl2 (g) + 2 Na (s) --------> Substance oxidized Oxidizing agent 2 Cl- (aq) 2 Na+ (aq) + Na Substance reduced Reducing agent Cl2 Cl2 Na WS # 3 Spontaneous and Non-spontaneous Redox Reactions Describe each reaction as spontaneous or non-spontaneous. 1. Au+3 Fe+3 -----> Fe+3 ------> Fe+2 F- ------> F2 + 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 + 2. Pb + 3. Cl2 + 4. S2O8-2 Fe+2 + Au Pb+2 + + + 2Cl- nonspontaneous (two oxidizing agents) spontaneous 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 Cl- O2 Na+ Fe 13. Determine the oxidation number for the element underlined. 14. 15. PbSO4 HP032- 6 3 ClO3Na2O2 5 -1 CaH2 NaIO3 -1 5 Al2(SO4)3 C4H12 6 -3 Al3+ + Zn ---------> Substance oxidized Zn Cr2O72- + ClO2- ------------> Substance reduced Al + Zn2+ Oxidizing agent Cr3+ Cr2O72- + Al3+ ClO4- Oxidizing agent Cr2O72- 16. State the Oxidation Number of each of the elements that is underlined. a) NH3 c) ZnCO3 e) Na -3 4 0 b) H2SO4 d) Al(OH)3 f) Cl2 6 3 0 17. Balance the redox equation using the half reaction method. Al + 3Ag+ ----------> Al3+ + 18. Circle each formula that is able to lose an electron O2 Cl- Na+ Fe Determine the oxidation number for the element underlined. 19. 20. PbSO4 ClO3- 2 5 3Ag 21. 22. HPO32Na202 3 -1 23. 24. 25. 26. CaH2 NaIO3 C4H12 Al2(SO4)3 -1 5 -3 6 27. Al3+ + Zn ----------> Substance oxidized Cr2O72- + 28. Substance oxidized Cr3+ Cr2O72- O3 + H2O + SO2 -----> SO42- + O2 ClO4- + 2H+ Reducing agent + 4NO3- + 7H2O + 4 H+ --------> 6H3AsO4 NO3- WS # 4 + Oxidizing agent SO2 Substance reduced Al3+ Oxidizing agent ClO2- ----------------> Cr2O72- 30. 3As2O3 Zn2+ + Zn Substance reduced 29. Al SO2 + 4NO Reducing agent As2O3 Balancing Redox Reactions Balance each of the following half-cell reactions. (In each case assume that the reaction takes place in an ACIDIC solution.) Also, state whether the reaction is oxidation or reduction. 1. 5H2O + S2O32- --------------> 2SO42- + 10H+ + 8e- oxidation 2. 8H+ + 5e- + MnO4- --------------> Mn2+ + 4H2O reduction 3. 4H2O + As --------------> AsO43- + 8H+ + 5e- Cr2O72- + 14H+ + 6e- oxidation 4. 7H2O + 2Cr3+ -----------> oxidation 5. Pb2+ 2H2O + --------------> PbO2 + 4H+ + 2e- oxidation 8H+ 6. + SO42- + 6e- --------------> S NO3- + 3e- -------------> + 4H2O reduction 4H+ 7. + NO + 2H2O NH4+ + 3H2O reduction 10H+ + 8. 8e- + NO3- --------------> 10e- + 2BrO3- --------------> reduction 12H+ + 9. Br2 + reduction Balancing Half Cell Reactions Balance in basic solution. 10. 3e- 11. 4H2O + 12. 8OH- 13. 14OH- 14. 4OH- 15. 4H2O + 16. 10 OH- 17. 7H2O + + 2H2O 5e- + As + + + NO3- --------------> NO MnO4- --------------> AsO43- + Pb2+ --------------> 6e- + 4H2O + 8e- + 8OH- 5e- + 7H2O + 6e- PbO2 + 2H2O + 2e- SO42- --------------> S S2O32- --------------> 2SO42- + + 4OH- --------------> Mn2+ + 2Cr3+ --------------> Cr2O72- + + + 5H2O + NO3- --------------> NH4+ + 8OH8e- 10 OH- 6H2O 18. 10e- 6H2O + + 2BrO3- --------------> Br2 + 12 OH- 19. Determine if each of the following changes is oxidation, reduction or neither. 20. SO32- --------> SO42- oxidation CaO --------> Ca reduction CrO42- --------> Cr2O72- neither CrO42- --------> Cr3+ reduction 2I- --------> I2 oxidation IO3- --------> I2 reduction MnO4- --------> Mn2+ reduction ClO2- --------> ClO- reduction Cr2O72- Fe2+ + --------> Cr3+ + Fe3+ Substance oxidized Fe2+ Substance reduced Cr2O72Oxidizing agent Cr2O72- Reducing agent Fe2+ WS #5 Balancing Redox Reactions in Acid and Basic Solution Balance each redox equation. Assume all are spontaneous. Use the half reaction method. 1. 2O2- + 2F2 -----------> O2 + 4F- 2. 4Al + 3O2 -----------> 6O2- + 4Al3+ 3. 2K + Zn+2 -----------> Zn + 2K+ Balance each half reaction in basic solution. 4. Cr2O72- + 5. NO 4OH- 6. 2H2O 7. 2MnO2 + + 7H2O + 6e- --------------> 14OH- ------------------> 2H2O --------------> SO2 --------------> Mn2O3 2e- + SO42- + H2 O + 2e- + 2Cr3+ + NO3- + 3e4OH- + + 2OH- Balance each redox reaction in acid solution using the half reaction method. 8. 8H+ + Cr2O72- 3H2O2 + -------> 3O2 2Cr3+ + + 7H2O 9. TeO32 - + 2N2O4 10. 4H+ + 4ReO4-+ 7IO- 11. 8H+ + 5PbO2 + I2 12. 12H2O + 8As + H2 O -------> 4NO3- + 2H+ -------> 7IO3- + 4Re + 2H2O -------> 5Pb2+ + 2IO3- + 4H2O 3H2AsO4- 5AsH3 + -------> Te + + 3H+ Balance each redox reaction in basic solution using the half reaction method. + 8OH- + 2Cr3+ 13. 3O2 14. H2 O 15. 7IO3- + 4OH- + 4Re 16. 8OH- + 5Pb2+ + 17. 7H2O + Te + -------> H2O 4NO3- -------> TeO32- + -------> 4ReO4- + + 2OH7IO- 2IO3- -------> 5PbO2 + + Cr2O72- + 3Hg -------> 3Hg2+ + 3H2O2 + + 2N2O4 + 2H2O Cr2O72- I2 + 4H2O 14OH- + 2Cr3+ State of the change represents oxidation, reduction or neither (use oxidation #s). 18. 19. 20. 21. 22. --------> MnO2 NH3 HClO4 O2 P2O5 Mn2O3 NO2 HCl + O2P4H10 --------> -------> --------> --------> reduction oxidation reduction reduction reduction H2O Determine the oxidation number 23. H2S O4 6 22. HS O4- 6 24. 25. P4 U O3 0 6 23. 24. NaH Na2O 2 -1 -1 26. U 2O5 5 25. P b SO4 2 WS #6 Review 1. Describe each in your own words 1. Oxidation 2. Reduction 3. Oxidizing agent 4. Reducing agent - loss of electrons - gain of electrons - causes oxidation by undergoing reduction - causes reduction by undergoing oxidation 2. Write half reactions for each. Describe as oxidation or reduction. Circle all oxidizing agents. a) b) c) d) e) f) Na Ca Al3+ 2F1N2 2O2- -----------> Na+ + e-----------> Ca2+ + 2e+ 3e -----------> Al ----------> F2 + 2e+ 6e ----------> 2N3----------> O2 + 4e- oxidation oxidation reduction oxidation reduction oxidation 3. Write the reaction between the following: Use the half reaction method. a) Ca + 3Ca b) Sn + Sn + Al(NO3)3 + 2Al3+ -------------> 2Al + AgNO3 2Ag+ -------------> 2Ag + Sn2+ 3Ca2+ c) Sn + 3Sn Au(NO3)3 + 2Au3+ -------------> 2Au + 3Sn2+ 4. Circle each reducing agent: Cu Cu+ Al Al3+ 5. Circle each oxidizing agent: F- F O2- O2 6. Ni+2 reacts with Mn, however, Al+3 does not react with Mn. Rank the oxidizing agents in order of decreasing strength. Rank the reducing agents in order of decreasing strength. strongest oxidizing agent Ni2+ + Mn2+ + Al3+ + 2e2e3e- -----------> -----------> -----------> Ni Mn 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+ Pb2+ Ca2+ + + + 1e2e2e- -----------> -----------> -----------> Ag Pb 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 + Ag+ + Mg2+ + 2e1e2e- --------> -----------> -----------> 2ClAg 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+ + Mn2+ + Al3+ + 2e2e3e- -----------> -----------> -----------> Ni Mn 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 Br2 I2 + + + 2e2e2e- --------> --------> --------> Classify as oxidation, reduction or neither. 11. SO4212. MnO2 13. Cr2O72- --------> --------> --------> S2MnO4CrO42- reduction oxidation neither 2Cl2Br2I- strongest reducing agent 14. IO3- --------> I2 reduction 15. Given the following lab data SnCl2 Ni(NO3)2 Cr(NO3)3 & & & Ni Fe Fe Spontaneous Spontaneous Non spontaneous. i) Write three balanced equations. Ni + Sn2+ -------------> Ni2+ Fe + Ni2+ -------------> Fe2+ Fe + Cr3+ <------------- Fe2+ + + + Sn Ni Cr ii) Rank the oxidizing agents in decreasing order of strength. strongest oxidizing agent Sn2+ Ni2+ Fe2+ Cr3+ + + + + 2e2e2e3e- -----------> -----------> -----------> -----------> Sn Ni Fe Cr strongest reducing agent iii) Rank the reducing agents in decreasing order of strength. See above. iv) Will SnCl2 react with Cr? Explain? Yes, because Sn2+ is a stronger oxidizing agent than Cr3+ . v) Will Fe2+ react with Sn? 2H+ 16. + No, because Fe2+ is a weaker oxidizing agent than Sn2+ 2MnO4- + 5H2S --------> oxidizing agent reducing agent 5S + 6H2O + 2MnO 17. 2H+ + 10SO42+ 4Br2 oxidizing agent 18. Balance in basic solution 2MnO419. 5H2S --------> 5S + 2MnO + 4H2O + 2OH- Describe as spontaneous or non-spontaneous. Use your reduction potential chart. a) b) c) d) 20. + ----------> 5S2O32- + 8BrO3- + H2O reducing agent ZnCl2 CuCl2 Br2 H2S & & & & Cu NaCl Fe2+ Al3+ nonspontaneous nonspontaneous spontaneous nonspontaneous Can you keep HCl in a Zn container? What about an Au container? No, Spontaneous reaction. Yes, nonspontaneous reaction. Balance in basic solution 21. H2 O + 10SO42- + 4Br2 ------> 5S2O32- + 2OH- + 8BrO3- Classify as an oxidizing agent, reducing agent or both based on its position on the table. State the Eoor voltage of its position. Some of these are both, so state two voltages and indicate that it can be an oxidizing and reducing agent. e.g. MnO4- (in acid) oxidizing agent 1.51 v 22. 23. 24. 25. 26. 27. oxidizing agent oxidizing agent / reducing agent oxidizing agent reducing agent oxidizing agent oxidizing agent / reducing agent 1.09 v -0.45 v / 0.77 v 0.60 v -0.26 v -0.74 v -0.40 v / +0.80 v Br2 Fe2+ MnO4- (water) Ni Cr3+ H2O Indicate as spontaneous or non-spontaneous. 28. 29. MnO4Cu2+ & & Fe2+ Br- non-spontaneous non-spontaneous 30. 31. 32. 33. HNO3 MnO4- (acid) Ni(s) HCl & & & & Ag H2O Al3+ Mg spontaneous spontaneous non-spontaneous spontaneous Write each oxidation and reduction half reaction for each question above. Determine the Eo for each. Calculate the Eo for the overall reaction. MnO4- 34. MnO4- + + 2H2O + 3e- --------> MnO2 3(Fe2+ -----------> Fe3+ + 1e-) 2H2O + 3Fe2+ -----------> 3Fe3+ + MnO2 + 4OH- + 4OH- +0.60 v -0.77 v -0.17 v 35. 36. NO3- + NO3- + 37. 4H+ +3e- -----------> NO + 3(Ag ----------> Ag + 1e-) 4H+ + 3Ag ----------> NO + 2H2O + 2H2O + 3Ag+ +0.96 v -0.80 v +0.16 v 38. 39. Mg 2H+ Mg + + 2e------> H2 ----------> Mg2+ + 2e+ 2+ 2H ----------> Mg + WS # 7 H2 0.00 v 2.37 v 2.37 v 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 . 2 e- 2 e- voltmeter 1.0 M KNO3 NO3- K+ Zn Cu Zn → Zn2+ + 2eoxidation anode 0.76 v loses mass Zn2+ Cu2+ NO3- NO3- 1 M Zn(NO3)2 1 M Cu(NO3)2 Cu2+ + Zn → Zn2+ + Cu 12. Ag / AgNO3 ║ H2 / HCl 1.10 v Cu has greater electron affinity Cu2+ + 2e- → Cu reduction cathode 0.34 v gains mass 2 e- 2 e- voltmeter 1.0 M KNO3 NO3- K+ H2 H2 → 2H+ + 2eoxidation anode 0.00 v Cu H+ Ag+ Cl- NO3- 1 M HCl 1 M Ag(NO3)2 2Ag+ + H2 → 2Ag + 2H+ 0.80 v Ag has a greater electron affinity 2Ag+ + 2e- → 2Ag reduction cathode 0.80 v gains mass WS # 8 1. In an electrolytic cell, reduction occurs at the negative electrode and oxidation occurs at the positive electrode. 2. If there are two possible reduction reactions, the highest one on the chart occurs. 3. For reduction, the chart is read from left to right. 4. For oxidation, the chart is read from right to left and the sign of the voltage is changed. 5. If there are two possible oxidation reactions, the lowest one on the chart occurs. 6. Corrosion of a metal is oxidation. 7. Electrolysis uses electrical energy. 8. Electrochemical cells produce electrical energy. 9. Electrolytic cells use electrical energy. 10. What is the standard reference cell? hydrogen Eo = O v Draw and completely analyze each electrolytic cell. 11. Molten NaCl Power Source - + Pt Pt Na+ Cl- Cathode: Na+ + 1e- → 2e-1.36 v Na(s) Anode: 2Cl- → Cl2 + -2.71 v Overall: 2Na+ + 2Cl- → Cl2 + 2Na(s) -4.07 v MTV = +4.07 v 12. Aqueous Na2SO4 Power Source - + C C Na+ SO42H2O Cathode: 2H2O + 2e- → H2 + 2OH- -0.41 v + 1/2O2 + 2e-0.82 v Anode: H2O Overall: H2O → H2 + 1/2O2 MTV = +1.23 v -1.23 v → 2H+ 13. Liquid K2O Power Source - + Pt Pt K+ O2- Cathode: K+ + 1e- → 4e?v K(s) Anode: 2O2- → O2 + -2.93 v Overall: 4K+ + 2O2- → O2 + 4K(s) -? v MTV = +? v 14. 1.0 M LiI Power Source - + Pt Pt Li+ I- Cathode: Cathode: 2H2O + 2e- → -0.54 v H2 + 2OH- Overall: 2H2O + 2I- → I2 + H2 + 2OH- -0.41 v -0.95 v Anode: 2I- → I2 + 2e- 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 + 2OH0.250L MnO4- x 0.200 mol x 2 mol MnO2 L 2 mol MnO4- x 86.9g = 4.34g mol 16. Determine the oxidation number for each underlined atom. MnO2 4 Cr2O7-2 6 IO35 C2O4-2 3 Al(NO3)3 5 17. Describe each term: Salt bridge- a u-tube filled with salt solution that allows ions to flow in an electrochemical cell. Electrolyte- a solution that conducts electricity Anode- an electrode that is the site of oxidation Cathode- an electrode that is the site of reduction Spontaneous- a reaction that occurs naturally and has a positive voltage Electron affinity- the ability of a metal to attract electrons 18. What would happen if you used an aluminum spoon to stir a solution of FeSO4(aq) ? Write a reaction and calculate Eo. 2Al + 3Fe2+ -------> 2Al3+ + would be a reaction! 3Fe E0 = 1.21 v Spontaneous. There 19. Draw an electrochemical cell using Cu and Ag electrodes. Cathode (+) Ag Ag+ + 1e---------> Ag 2Ag+ + Cu ------> 2Ag + 0.80v Anode (-) Cu Cu -------> Cu2 + 2e Cu2+ E0 = 0.46 v spontaneous -0.34v 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 2MnO40.100L = 1.88M 21. How is the breathalyzer reaction used to determine blood alcohol content (you might need to look this up in your textbook)? The breathalyzer reaction uses a spontaneous redox reaction between acidic Cr2O72- and ethanol C2H5OH. If alcohol is present in your breath sample, it will react with a solution of Cr2O72- reducing the orange color as it reacts to form Cr3+, which is green. The drunker you are, the greater the reduction in orange color, which is measured with a spectrophotometer. 22. 2H+ + Mg-----> Mg+2 +H2 H+ Oxidizing agent WS #9 Reducing agent Mg 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: Anode reaction: 1e- -------> Ag Overall reaction: Pb Pb Cathode: Cathode reaction: --------> Pb2+ + 2e- Pb + 2Ag+ -----> Pb2+ + 2Ag Ag Ag+ + Voltage: 0.93v 2. ZnCl2(l) electrolytic cell (electro-winning) Anode: Anode reaction: ---> Zn Overall reaction: C 2Cl- --------> Cl2 + 2e2Cl- + Zn2+ -----> Cl2 Cathode: Cathode reaction: + Zn Zn2+ C + 2e- ---- MTV: +2.12 v 3. CuSO4(aq) electrolytic cell (electro-winning) Anode: C Cathode: C Anode reaction: H2O --------> 2H+ Cu2+ + 2e- -------> Cu Overall reaction: H2O + Cu2+ + 1/2O2 + 2e-----> 2H+ Cathode reaction: + 1/2O2 + Cu MTV: +0.48 v 4. The electrolysis of 1M NaI (electro-winning) Anode: Anode reaction: ----> H2 + 2OHOverall reaction: C 2I- --------> I2 + 2e2H2O + 2I- Cathode: C Cathode reaction: 2H2O + 2e- --- -----> 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: Anode reaction: -------> Al Overall reaction: 6. + 4Al3+ -----> 3O2 + 4Al Ag Cathode: Ag-----> Ag+ + e- penny Cathode reaction: Ag+ + e- -----> penny Cathode reaction: Ni2+ + 2e- ----- The reaction needed to nickel plate a copper penny. Anode: Anode reaction: > Ni Possible Electrolyte 8. 6O2- Cathode: C Cathode reaction: Al3+ + 3e- 4e- + The reaction needed to electroplate a copper penny with silver. Anode: Anode reaction: Ag 7. C 2O2- -------> O2 Ni Cathode: Ni-----> Ni+2 + 2eNi(NO3)2 The reaction used in the electrorefining of lead. Anode: Anode reaction: Impure Lead Pb-----> Pb+2 + 2e- Cathode: Cathode reaction: Pure Lead 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 Anode reaction: Mg --------> Mg2+ + 2e+2 Zn + 2e -------> Zn Overall reaction: Mg + Zn2+ -----> Mg2+ + 2. Zn Voltage: 1.61v The electrolytic cell used to produce Al. Electrolyte: Anode: Anode reaction: -------> Al Overall reaction: 3. Cathode: Zn Cathode reaction: Al2O3 Phase (aqueous or molten) C 2O2- -------> O2 + 4e6O2- + 4Al3+ Molten Cathode: C Cathode reaction: Al3+ + 3e- -----> 3O2 + 4Al The electrolysis KI(aq) Anode: C Cathode: C Anode reaction: 2I- --------> I2 + 2eCathode reaction: 2H2O + 2e- -------> H2 + 2OHOverall reaction: 2H2O + 2I-----> H2 + 2OH- + I2 MTV: +0.95 v 4. The electrorefining of Pb Anode: Anode reaction: 5. Impure Lead Pb-----> Pb+2 + 2e- Cathode: Cathode reaction: Pure Lead Pb2+ + 2e- -----> Pb Nickel plating an iron nail. Anode: Anode reaction: > Ni Possible Electrolyte Ni Cathode: +2 Ni-----> Ni + 2eNi(NO3)2 nail Cathode reaction: Ni2+ + 2e- ----- The -ve side of the power supply is connected to the nail 6. Draw an Ag/ Zn electrochemical cell. Anode: Anode reaction: 1e- -------> Ag Overall reaction: Zn Zn --------> Zn2+ + 2e- Zn + 2Ag+ -----> Zn2+ + Cathode: Cathode reaction: 2Ag Ag Ag+ + Voltage: 1.56v 7. Draw a KF(l) electrolytic cell. Anode: Anode reaction: Overall reaction: C Cathode: C + 2F --------> F2 + 2e Cathode reaction: K + e -------> K 2F- + 2K+ -----> Cl2 + K MTV: +5.80v 8. Draw a KF(aq) electrolytic cell. Anode: C Anode reaction: H2O --------> 2H+ + 1/2O2 + 2e2H2O + 2e -------> H2 + 2OHOverall reaction: H2 O -----> H2 + 1/2O2 Cathode: Cathode reaction: C MTV: +1.23 v 9. Draw a FeI2(aq) electrolytic cell. Anode: Anode reaction: 2e- -------> Fe Overall reaction: C 2I- --------> I2 + 2eFe2+ + 2I- Cathode: C 2+ Cathode reaction: Fe + -----> 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 Anode reaction: Pb --------> Pb2+ + 2e+2 Cd + 2e -------> Zn x volts Overall reaction: Pb + Cd2+ -----> Pb2+ + 0.13 + x = 0.50 0.13v Cd Cathode: Cathode reaction: Cd Voltage: 0.50v x = 0.37v 11. Write the overall reaction and describe the anode and cathode for a dry (Leclanche), fuel, alkaline and lead/acid cell. Cell anode anode reaction cathode cathode reaction Leclanche or Common Dry Cell Zn Zn-->Zn+2 + 2e- C Mn+4 +1e- -----> Mn+3 NH4Cl and MnO2 Alkaline Cell Zn Zn-->Zn+2 + 2e- C Mn+4 +1e- -----> Mn+3 KOH and MnO2 Lead Storage or Pb Car Battery Pb ---> Pb+2+ 2e- PbO2 PbO2 + HSO4- + 3H+ + 2e- H2SO4 ----> PbSO4 + 2H2O Fuel Cell H2 + 2OH- ---> 2H2O + 2e- ½O2 + H2O +2e-----> 2OH- KOH C C electrolyte 12. 2HIO3 + 5H2S O3 ----------> I2 + 5H2SO4 + H2O oxidizing agent HIO3 substance reduced HIO3 substance oxidized reducing agent H2SO3 H2SO3 13.What is the electrolyte in a fuel cell? KOH 14. What is the fuel in a fuel cell? H2 and O2 15. Describe the differences and similarities between an electrolytic and electrochemical cell. Electrolytic Electrochemical Uses electricity Nonspontaneous Makes chemicals Inert carbon electrodes The negative electrode is reduction Produces electricity Spontaneous Uses chemicals Usually has a salt bridge The higher metal is reduction Oxidation occurs at the anode and reduction occurs at the cathode. Anions migrate to the anode and cations migrate to the cathode. Electrons go from anode to cathode through the wire. 16. Describe and give two examples of electrowinning. The electrolysis of water to make H2 and O2. The electrolysis of Al2O3 to make Al and O2. 17. Describe and give one example of electrorefinning. The electrorefinning of Pb. 18. List three metals that can be won from aqueous solution. Cu Fe Sn 19. List three metals that cannot be won from aqueous solution. Al Na Pb Au Ag Zn K Li Ca Mg 20. What is the electrolyte in a fuel cell, alkaline battery, Dry Cell (Leclanche) and lead acid battery? KOH KOH & MnO2 NH4Cl & MnO2 PbSO4 21. State two metals that can be used to cathodically protect Fe. Describe how they protect iron from corrosion. Zn and Mg. When attached to Fe they form an electrochemical cell. Zn or Mg is a stronger reducing agent (lower on the chart) and is the anode and Fe is the cathode. Since the cathode is the site of reduction, Fe cannot oxidize or corrode. 22. Write the half reaction that describes the corrosion of iron. Fe --------> Fe2+ +2e- 23. Write the half reaction that describes the reduction reaction that occurs when iron corrodes in air and water. 2e- + H2O + 1/2O2 ----------> 2OH- 24. Why does iron corrode faster in salt water? the rate of reaction in an electrochemical cell. The salt acts like a salt-bridge and increases 25. Write the anode and cathode reaction in an electrolytic cell with a CaCl2 (l) electrolyte. Cathode: Ca2+ + 2e- ---------> Ca Anode: 2Cl- ----------> Cl2 + 2e- 26. Explain why you would choose Zn or Cu to cathodically protect iron? Zn. It is a stronger reducing agent than Fe and it will allow Fe to be the cathode, which cannot corrode. 27. Choose a suitable redox reactant to oxidize Cl- to ClO4- in a redox titration. MnO4- in acid gives a spontaneous reaction as well as a color change from purple to clear. 28. Describe as an electrochemical or electrolytic cell: a) Fuel cell electrolytic c) Discharging a car battery electrolytic e) Industrial Al production electrolytic electrochemical b)Charging a car battery electrochemical d) Ni plating electrolytic f) Cl2 production 29) Write the anode and cathode reactions. Cell anode anode reaction cathode cathode reaction electrolyte Cl2 production C 2Cl- ------> Cl2 C + 2e- Na+ + e- -----> Na NaCl(l) Leclanche or Common Dry Cell Zn Zn-->Zn+2 + 2e- C/MnO2 Mn+4 +1e- -----> Mn+3 NH4Cl and MnO2 Nickel Plating Ni Ni-->Ni+2 + 2e- Metal to be Ni2+ +2e- -----> Ni plated Lead Storage or Car Battery Pb Pb ---> Pb+2+ 2e- PbO2 PbO2 + SO4-2 + 4OH-1 + 2e-----> PbSO4 + 2H2O H2SO4 Fuel Cell C H2 + 2OH- ---> 2H2O + 2e- C O2 + 2H2O +4e-----> 4OH- KOH Ni(NO3)2 30) Al and AgNO3(aq) are mixed and the surface of the Al darkens. List the two oxidizing agents in decreasing strength. List the two reducing agents in decreasing strength. Oxidizing Agents Ag+ Al3+ Reducing Agents Al Ag
