BALANCING REDOX EQUATIONS EXERCISE

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1 Ag + NO3 Ag 1+ + NO N2H4 + H2O2 N2 + H2O CO + Fe2O3 FeO + CO NO3 + CO CO2 + NO H2 + Fe3O4 Fe + H2O 0 +8/ BALANCING REDOX EQUATIONS EXERCISE H2C2O4 + MnO4 CO2 + MnO Zn + NO3 Zn 2+ + NO C2N2 CN + CNO ClO2 + SbO2 ClO2 + Sb(OH) (1) Cr2O7 2 + I Cr 3+ + I

2 Fe3O4 + H2O2 Fe 3+ + H2O +8/ MnO4 + NH3 MnO2 + NO CN + CrO4 2 CNO + Cr(OH) (1) Fe(CN)6 3 + Cr2O3 Fe(CN)6 4 + CrO (1) (1) NH4NO3 N2O NO2 + MnO4 NO3 + Mn2+ (in acid solution) I + MnO4 I2 + MnO2 (in basic solution) Cl2 + S2O3 2 Cl + SO4 2 (in acidic solution) CH4 + O2 C + H2O Br2 Br + BrO3 (in basic solution)

3 BALANCING REDOX EQUATIONS EXERCISE 1. Ag + NO3 Ag 1+ + NO ox: 3 (Ag Ag e ) 0 +1 red: NO3 + 3 e NO Ag + NO3 3 Ag 1+ + NO 3 Ag + NO3 + 4 H + 3 Ag 1+ + NO + 2 H2O 2. N2H4 + H2O2 N2 + H2O ox: N2H4 N2 + 4 e 2 0 red: 2 (H2O2 + 2 e 2 H2O) 1 2 N2H4 + 2 H2O2 N2 + 4 H2O 3. CO + Fe2O3 FeO + CO ox: CO CO2 + 2 e 2 +4 red: Fe2O3 + 2 e 2 FeO CO + Fe2O3 2 FeO + CO2 4. NO3 + CO CO2 + NO ox: CO CO2 + 2 e red: 2 (NO3 + 1 e NO2) NO3 + CO CO2 + 2 NO2 2 NO3 + CO + 2 H + CO2 + 2 NO2 + H2O

4 5. H2 + Fe3O4 Fe + H2O 0 +8/ ox: 4 (H2 H2O + 2 e ) 0 +1 red: Fe3O4 + 8 e 3 Fe +8/3 0 4 H2 + Fe3O4 3 Fe + 4 H2O 6. H2C2O4 + MnO4 CO2 + MnO ox: 5 (H2C2O4 2 CO2 + 2 e ) red: 2 (MnO4 + 5 e MnO) H2C2O4 + 2 MnO4 10 CO2 + 2 MnO 5 H2C2O4 + 2 MnO4 + 2 H + 10 CO2 + 2 MnO + 6 H2O 7. Zn + NO3 Zn 2+ + NO ox: 3 (Zn Zn e ) 0 +2 red: 2 (NO3 + 3 e NO) Zn + 2 NO3 3 Zn NO 3 Zn + 2 NO3 + 8 H + 3 Zn NO + 4 H2O 8. C2N2 CN + CNO ox: C2N2 2 CNO + 2e red: C2N2 + 2e 2 CN C2N2 2 CN + 2 CNO C2N2 + H2O CN + CNO + 2 H +

5 9. ClO2 + SbO2 ClO2 + Sb(OH) (1) ox: SbO2 Sb(OH)6 + 2e +3 red: 2 (ClO2 + 1 e ClO2 ) ClO2 + SbO2 2 ClO2 + Sb(OH)6 2 ClO2 + SbO2 + 2 OH + 2 H2O 2 ClO2 + Sb(OH)6 10. Cr2O7 2 + I Cr 3+ + I ox: 3 (2 I I2 + 2e ) 1 0 red: Cr2O e 2 Cr Cr2O I 2 Cr I2 Cr2O I + 14 H + 2 Cr I2 + 7 H2O 11. Fe3O4 + H2O2 Fe 3+ + H2O +8/ ox: 2 (Fe3O4 3 Fe e ) +8/3 +3 red: H2O2 + 2 e 2 H2O Fe3O4 + H2O2 6 Fe H2O 2 Fe3O4 + H2O H + 6 Fe H2O 12. MnO4 + NH3 MnO2 + NO ox: 3 (NH3 NO3 + 8e ) 3 +5 red: 8 ( MnO4 + 3 e MnO2) MnO4 + 3 NH3 8 MnO2 + 3 NO3 8 MnO4 + 3 NH3 + 5 H + 8 MnO2 + 3 NO3 + 7 H 2 O

6 13. CN + CrO4 2 CNO + Cr(OH) (1) ox: 3 (CN CNO + 2 e ) red: 2 (CrO e Cr(OH)3) CN + 2 CrO H 2 O 3 CNO + 2 Cr(OH)3 + 4 OH 3 CN + 2 CrO H 2 O 3 CNO + 2 Cr(OH)3 + 4 OH 14. Fe(CN)6 3 + Cr2O3 Fe(CN)6 4 + CrO (1) (1) +6 2 ox: Cr2O3 2 CrO e red: 6 (Fe(CN) e Fe(CN)6 4 ) Fe(CN)6 3 + Cr2O3 + 5 H 2 O 6 Fe(CN) CrO H + 6 Fe(CN)6 3 + Cr2O3 + 5 H 2 O 6 Fe(CN) CrO H NH4NO3 N2O ox: 2 NH4 + N2O + 8e 3 +1 red: 2 NO3 + 8 e N2O NH NO3 2 N2O NH4NO3 N2O + 2 H2O 16. NO 2 + MnO 4 NO3 + Mn 2+ (in acid solution) ox: 5 (NO 2 NO3 + 2 e ) red: 2 (MnO e Mn 2+) NO MnO H + 5 NO Mn H 2 O 5 NO MnO H + 5 NO Mn H 2 O

7 17. I + MnO 4 I 2 + MnO 2 (in basic solution) ox: 3 (2 I I2 + 2 e ) 1 0 red: 2 (MnO e MnO 2 ) I + 2 MnO 4 3 I MnO 2 6 I + 2 MnO H 2 O 3 I MnO OH 18. Cl2 + S2O3 2 Cl + SO4 2 (in acidic solution) ox: S2O3 2 2 SO e red: 4 (Cl2 + 2 e 2 Cl ) Cl2 + S2O H2O 8 Cl + 2 SO H + 4 Cl2 + S2O H2O 8 Cl + 2 SO H CH4 + O2 C + H2O ox: CH4 C + 4 e 4 0 red: O e 2 H 2 O 0 2 CH4 + O2 CH4 + O2 C C + 2 H2O 20. Br2 Br + BrO3 (in basic solution) ox: Br2 2 BrO e 0 +5 red: 5 (Br2 + 2 e 2 Br ) Br2 10 Br + 2 BrO3 3 Br2 + 6 OH 5 Br + BrO3 + 3 H2O

8 Predicting REDOX Reactions Building a REDOX Table 1. The following reactions were performed. Construct a table of relative strengths of oxidizing and reducing agents written as reductions and with the SOA to WOA. Zn + Co 2+ Zn 2+ + Co Mg 2+ + Zn no rxn Half rxns: Final Order: Zn e Zn SOA Co e Co X Co e Co Zn e Zn Mg e Mg Mg e Mg SRA 2. In a school laboratory four metals were combined with each of four solutions. Construct a table of relative strengths of oxidizing and reducing agents written as reductions and with the SOA to WOA. Be + Cd 2+ Be 2+ + Cd Cd + 2 H + Cd 2+ + H 2 Ca 2+ + Be no rxn Cu + 2 H + no rxn Half rxns: Final Order: Be e Be SOA Cu e Cu X Cd e Cd 2 H + + 2e H 2 X 2 H+ + 2e H 2 Cd2+ + 2e Cd Ca e Ca Be e Be X Cu e Cu Ca e Ca SRA 3. Write and rank the two half reaction equations for each of the following reactions: (a) Co + Cu(NO 3 ) 2 Cu + Co(NO 3 ) 2 Cu e Cu Co e Co

9 (b) Cd + Zn(NO 3 ) 2 Zn + Cd(NO 3 ) 2 Zn e Zn Cd e Cd (c) Br 2 + 2KI I KBr Br 2 + 2e 2 Br I 2 + 2e 2 I 4. Prepare a REDOX table of halfreactions showing the relative strengths of oxidizing and reducing agent for the following: OA Al 3+ Tl + Ga 2+ In 3+ RA Al X Tl X X X X Ga X X In X X X WOA Rank 4 th SOA Rank 1 st Rank 3 rd Rank 2 nd SOA Tl + + e Tl In e In Ga e Ga Al e Al SRA Prediction REDOX Reaction in Solution 1. List all the entities initially present in the following mixtures and identify all possible oxidizing and reducing agents. Write the resulting REDOX reaction (or no rxn). (a) A lead strip is placed in a copper (II) sulfate solution. OA (Cu +, Cu) (not in H + ) (H2) Pb SRA Cu 2+ SOA SO4 H2O RA (Pb 2+ ) (S2O8 2 ) (O2) ox: Pb Pb e red: Cu e Cu Pb + Cu 2+ Pb 2+ + Cu

10 (b) A potassium dichromate solution is added to an acidic iron (II) nitrate solution. OA (K) (Cr 3+ ) (Fe) (NO2) (H2) K + Cr2O7 2 SOA Fe 2+ SRA H +, NO3 H2O RA (Fe 3+ ) (O2) 6x ox: 6 Fe 2+ 6 Fe e red: Cr2O e 2 Cr 3+ Cr2O Fe H + 2 Cr Fe H2O (c) An aqueous chlorine solution is added to a phosphorous acid solution. OA (Cl ) RA (H2) Cl2 SOA H +, PO3 3 H2O SRA (O2) ox: 2 H2O O2 + 4 H e 2x red: 2 Cl2 + 4 e 4 Cl 2 Cl2 + 2 H2O 4 Cl + O2 + 4 H + (d) A potassium permanganate solution is mixed with an acidified tin (II) chloride solution. OA (K) (Mn 2+ ) (Sn) (H2) K + MnO4, H + SOA Sn 2+ SRA Cl H2O RA (Sn 4+ ) (Cl2) (O2) 5x ox: 5 Sn 2+ 5 Sn e 2x red: 2 MnO e 2 Mn 2+ 2 MnO4 + 5 Sn H + 2 Mn Sn H2O

11 Electrochemical (Galvanic or Voltaic) Cells Worksheet 1. a) Determine the anode, cathode and calculate the standard cell potential produced by a galvanic cell consisting of a Ni electrode in contact with a solution of Ni 2+ ions and a Ag electrode in contact with a solution of Ag 1+ ions. Ni 2+ (aq) + 2e Ni (s) E = 0.26 V (lesser flip) Ag + (aq) + e Ag (s) E = V ANODE: Ni (s) Ni 2+ (aq) + 2e E = V CATHODE: 2 Ag + (aq) + 2e 2 Ag (s) b) Write the shorthand cell notation. Ni (s) Ni 2+ (aq) Ag 1+ (aq) Ag (s) E = V E = V 2. a) Determine the anode, cathode and calculate the voltage produced by a galvanic cell consisting of an Fe electrode in contact with a solution of Fe 2+ ions and a Al electrode in contact with a solution of Al 3+ ions. Fe 2+ (aq) + 2e Fe (s) Al 3+ (aq) + 3e Al (s) E = 0.44 V E = 1.66 V (lesser flip) ANODE: 2 Al (s) 2 Al 3+ (aq) + 6e E = V CATHODE: 3 Fe 2+ (aq) + 6e 3 Fe (s) b) Write the shorthand cell notation. Al (s) Al 3+ (aq) Fe 2+ (aq) Fe (s) E = 0.44 V E = V 3. a) Determine the anode, cathode and calculate standard cell potential produced by a galvanic cell consisting of a C electrode in contact with an acidic solution of ClO 4 ions and a Cu electrode in contact with a solution of Cu 2+ ions. Which is anode and which is the cathode? ClO 4 (aq) + 8H + (aq) + 8e Cl (aq) + 4H 2 O (l) Cu 2+ (aq) + 2e Cu (s) E = V E = V (lesser flip) ANODE: 4 Cu (s 8 Cu 2+ (aq) ) + 8e E = 0.34 V CATHODE: ClO 4 (aq) + 8H + (aq) + 8e Cl (aq) + 4H 2 O (l) b) Write the shorthand cell notation. Cu (s) Cu 2+ (aq) ClO 4, H + (aq) C (s) E = V E = V

12 4. An electrochemical cell is constructed using electrodes based on the following half reactions: Pb 2+ (aq) + 2e Pb (s) Au 3+ (aq) + 3e Au (s) a) Which is the anode and which is the cathode in this cell? ANODE: Pb CATHODE: Au b) What is the standard cell potential? ANODE: 3 Pb (s) 3 Pb 2+ (aq) + 6e E = V CATHODE: 2 Au 3+ (aq) + 6e 2 Au (s) E = V E = V 5. Use complete halfreactions and potentials to predict whether the following reactions are spontaneous or nonspontaneous in aqueous solutions. If the cell is spontaneous, write the cell shorthand notation. a) Ca 2+ (aq) + 2 I (aq) Ca (s) + I 2(aq) ANODE: 2 I (aq) I 2(aq) + 2e E = 0.54 V CATHODE: Ca 2+ (aq) + 2e Ca (s) E = 2.87 V E = 3.41 V E is negative, therefore the cell is nonspontaneous. b) 2 H 2 S (g) + O 2(g) 2 H 2 O (l) + 2 S (s) ANODE: H 2 S (g) 2 S (s) + 2H + (aq) + 2e E = 0.14 V CATHODE: O 2(g) + 4H + (aq) + 4e 2 H 2 O (l) E = V E = V E is positive, therefore the cell is spontaneous. Pt (s) H 2 S (g) ; S (g) O 2 (g), H + (aq) C (s) c) SO 2(g) + MnO 2(s) Mn 2+ (aq) + SO 4 2 (aq) ANODE: SO 2(g) + 2 H 2 O (l) SO 4 2 (aq) + 4H + (aq) + 2e E = 0.18 V CATHODE: MnO 2(s) + 4H + (aq) + 2e Mn 2+ (aq) + 2 H 2 O (l) E is positive, therefore the cell is spontaneous. Pt (s) SO 2(g) ; SO 4 2 (aq) MnO 2(s), H + (aq) ; Mn2+ (aq) C (s) E = V E = V d) 2 H + (aq) + 2 Br (aq) H 2(g) + Br 2(aq) ANODE: 2 Br (aq) Br 2(l) + 2e E = 1.07 V CATHODE: 2 H + (aq) + 2e H 2(g) E = 0.00 V E = 1.07 V E is negative, therefore the cell is nonspontaneous.

13 e) Ce 4+ (aq) + Fe2+ (aq) Ce3+ (aq) + Fe3+ (aq) ANODE: Fe 2+ (aq) Fe3+ (aq) + e E = 0.77 V CATHODE: Ce 4+ (aq) + e Ce 3+ (aq) E = V E = 2.38 V E is negative, therefore the cell is nonspontaneous. f) Cr 2+ (aq) + Cu2+ (aq) Cr3+ (aq) + Cu+ (aq) ANODE: Cr 2+ (aq) Cr3+ (aq) + e E = V CATHODE: Cu 2+ (aq) + e Cu + (aq) E = V E = V E is positive, therefore the cell is spontaneous. C (s) Cr 2+ (aq) ; Cr 3+ (aq) Cu 2+ (aq) ; Cu + (aq) C (s)

14 Electrolytic Cells Worksheet 1. a) Give the cathode, anode and overall equations including cell potentials to conclude what happens to the ph of the solution near the cathode and anode during the electrolysis of KNO 3? Consider all possible reactions. OA K (2.92) H 2 (0.83) K + NO 3 SRA H 2 O SOA RA O 2 (1.23) ox: 2 H 2 O O 2 + 4H e E ox = 1.23 V 2x red: 2 H 2 O + 4 e 2H 2 + 2OH E red = 0.83 V 6 H 2 O O 2 + 2H 2 + 4H OH E cell = 2.06 V 2 H 2 O O 2 + 2H 2 E cell = 2.06 V at the anode ph, at the cathode ph b) Write the shorthand cell notation. C(s) C(s) or Pt(s) KNO 3(aq) or Pt(s) 2. Given the following molten systems, predict the products at each electrode. Assume inert electrodes and sufficient voltage to cause a reaction to take place. Consider all possible rxns. a) FeBr 2 OA Fe (0.45) SRA Fe 2+ SOA Br RA Fe 3+ (0.77) Br 2 (1.07) Fe 3+ is produced at the anode, Fe at the cathode. b) NiCl 2 OA RA Ni (0.26) Ni 2+ SOA Cl SRA Cl 2 (1.36) Cl 2 is produced at the anode, Ni at the cathode.

15 c) Na 2 SO 4 OA Na (2.71) Na + SOA SO 4 2 SRA RA S 2 O 4 2 (2.01) S 2 O 4 2 is produced at the anode, Na at the cathode. 3. Given the following 1.00 M solutions at 25 C predict the anode and cathode half cell reactions. What is the minimum voltage required for each cell to operate? a) LiMnO 4 OA Li (3.00) MnO 2 (+0.60) H 2 (0.83) Li + MnO 4 SOA H 2 O SRA RA O 2 (1.23) E cell = V = 0.63 V ; 0.63 V are needed b) CrI 3 OA Cr (0.76), Cr 2+ (0.41) H 2 (0.83) Cr 3+ SOA I SRA H 2 O SRA RA I 2 (0.54) O 2 (1.23) E cell = V = 0.95 V ; 0.95 V are needed c) Sn(NO 3 ) 2 OA Sn (0.14) H 2 (0.83) SRA Sn2+ SOA NO 3 H 2 O SRA RA Sn 4+ (0.15) O 2 (1.23) E cell = V = 0.29 V ; 0.29 V are needed d) Ag 2 SO 4 OA Ag (+0.80) SO 2 3 (0.93) H2 (0.83) Ag + SOA SO 2 4 H 2 O SRA RA S 2 O 2 8 (2.01) O2 (1.23) E cell = V = 0.43 V ; 0.43 V are needed

16 Stoichiometry and Free Energy Worksheet 1. How many coulombs, q, are required to deposit g of Ni from a solution of Ni 2+? Ni e Ni m = g M = g/mol n q q F then q n e e 1mol Ni g x g 1930C 1.93x 10 x F 3 x 1mol Cu C 2 mol e 2 x 9.65 x 10 mol e 4 C 2. Three electrolysis cells are connected in series. They contain, respectively, solutions of copper (II) nitrate, silver nitrate, and chromium (III) sulfate. If 1.00 g of copper is electrochemically deposited in the first cell, calculate the mass of silver and chromium deposited in the other cells. 1) Cu e Cu m = 1.00 g M = g/mol n e (Cu) 1.00g x 1mol Cu 63.55g 2 mol e x 1mol Cu mol 2) Ag + + e Ag n = mol m =? M = g/mol m Ag mol e 1mol Ag x 1mol e g x 1mol Ag 3.40g 3) Cr e Cr n = mol m =? M = g/mol mcr 1mol Cr mol e x 3 mol e g x 1mol Cr g

17 3. A constant current of 3.7 milliampere is passed through molten sodium chloride for 9.0 minutes. The sodium produced is allowed to react with water (500 ml). What is the ph of the resulting solution? Na + + e Na I = 3.7 ma n =? t = 9.0 min = 540 s n Na 3.7 x 10 s 2.1x C x 540s x mol mol e 9.65x 10 4 C 1mol Na x 1mol e 2 Na + 2 H 2 O 2 NaOH + H 2 n = 2.1 x 10 5 mol C =? V = 0.500L C NaOH ph 14 log (4.2 x x x mol NaOH mol x x 2 mol Na M 5 ) L 4. Given these halfreactions and their standard reduction potentials. 2 ClO H e Cl H 2 O E o (ClO 4 ) = V S 2 O e 2 SO 4 2 E o (S 2 O 8 2 ) = V Calculate: (a) Complete the REDOX reaction. an (ox): Cl H 2 O 2 ClO H e E o = 1.47 V 5 x cat (red): 5 S 2 O e 10 SO 4 2 E o = V 5 S 2 O Cl H 2 O 10 SO ClO H + (b) E o cell = V = V (b) ΔG o for the cell reaction G o = n F E o cell 10 mol e 9.65 x 10 4 C 0.54 J x x J 520 kj mol e C

18 (c) K eq for the cell reaction. ln Keq Keq J mol RT G J mol 1 K 1 x 298 K ( ) e RT G e ((210)) 1.6 x The system 2 AgI + Sn Sn Ag + I has a calculated E o cell = V. What is the value of K eq for this system? ln Keq Keq nfe o RT cell ( e nfe RT cell 2 mol e x 9.65 x 10 4 C x J C J mol 1 K 1 x 298 K ) e (1.2) Calculate ΔG o for the following reaction, given that its standard cell potential is V at 25 o C. NiO Cl + 4 H + Cl 2 + Ni H 2 O G o = n F E o cell 2 mol e 9.65 x 10 4 C J x x J 67.6 kj mol e C

19 Review Questions for SCH 4U Electrochemistry Test 1. Balance the following REDOX reaction in acidic solution Zn + NO 3 Zn 2+ + NH ox: 4 (Zn Zn e ) 0 +2 red: NO3 + 8 e NH Zn + NO 3 4 Zn + NO 3 4 Zn 2+ + NH H + 4 Zn 2+ + NH H 2 O 2. Balance the following REDOX reaction in acidic solution MnO4 + C2 O 4 2 CO2 + MnO ox: 3 (C 2 O CO2 + 2 e ) red: 2 (MnO e MnO 2 ) MnO 4 2 MnO C 2 O C 2 O MnO CO H + 2 MnO CO H 2 O 3. Given the following reactions, generate a standard reduction potential table: W 2+ + Z Z 2+ + W X 2+ + W W 2+ + X X 2+ + Y no rxn SOA Y e Y X e X W e W Z e Z SRA 4. Describe and explain what will happen if carbon electrodes are placed in a FeCl 2 solution. Give ALL possible half reactions. Species in Solution: Fe 2+ Cl 1 H 2 O Possible Reduction Half Reactions (OA): Fe e Fe E o = 0.44 V SOA 2 H 2 O + 2e O 2 + 4H + + 4e E o = 0.83 V

20 Possible Oxidation Half Reactions (RA): 2 Cl 1 Cl 2 + 2e E o = 1.36 V Fe 2+ Fe 3+ + e E o = 0.77 V SRA 2 H 2 O O 2 + 4H + + 4e E o = 1.23 V Full Reaction: 3 Fe 2+ Fe + 2 Fe 3+ E o cell = 1.21 V 5. Use the redox spontaneity rule to predict whether the following mixtures will be spontaneous or not. (a) Nickel metal in a solution of silver ions OA Ag (+0.80) H 2 (0.83) Ni SRA Ag + SOA H 2 O RA Ni 2+ (+0.26) O 2 (1.23) E o cell = V V = V spontaneous (b) Chlorine gas bubbled into a bromide ion solution OA Cl (+1.36) H 2 (0.83) Cl 2 SOA Br SRA H 2 O RA Br 2 (1.07) O 2 (1.23) E o cell = V 1.07 V = V spontaneous (c) Copper metal in nitric acid OA NO 2 (+0.80) H 2 (0.83) Cu SRA NO 3, H + SRA H 2 O RA Cu + (0.52) (1.07) O 2 (1.23) Cu 2+ (0.34) E o cell = V 0.34 V = V spontaneous 6. Three electrolysis cells are connected in series. They contain, respectively, solutions of zinc nitrate, aluminum nitrate and silver nitrate. If 1.00 g of silver is deposited in the third cell what mass of aluminum and zinc were deposited in the other cells. 1) Ag + + e Ag m = 1.00 g M = g/mol

21 ne 1.00 g x 1mol Ag g 1mol e x 1mol Ag 9.27 x 10 3 mol 2) Zn e Zn n = 9.27 x 10 3 mol m =? M = g/mol mzn 1mol Zn g 9.27 x 10 3 mol x x g 2 mol e 1mol Zn 3) Al e Al n = 9.27 x 10 3 mol m =? M = g/mol mal 7. For the cell: 1mol Al g 9.27 x 10 3 mol x x g 3 mol e 1mol Zn Ag (s) Ag1+ (aq) Zn2+ (aq) C (s) a) List all possible halfreactions that will occur at the cathode, including their cell potentials. Zn e Zn E o = 0.76 V SOA 2 H2O + 2 e 2 OH 1 + H2 E o = 0.83 V b) List the possible halfreactions that will occur at the anode, including their cell potentials. Ag Ag 1+ + e E o = 0.80 V SRA 2 H 2 O O 2 + 4H + + 4e E o = 1.23 V c) Give the full balanced REDOX reaction with the value for the cell s E o Zn 2+ + Ag Zn + Ag 1+ d) Draw a fully labeled diagram of the electrolytic cell. E o cell = 0.80 V (+0.76V) = 1.56 V

22 8. For the cell: Ag (s) S2 (aq) HCl (aq) Pt (s) a) List all the possible anode reactions with their E o values. 2 Ag + S 2 Ag 2 S E o = V SRA 2 H 2 O O 2 + 4H + + 4e E o = 1.23 V b) List all the possible cathode reactions with their E o values. 2 H2O + 2 e 2 OH 1 + H2 E o = 0.83 V 2 H + + 2e H 2 E o = 0.00 V SOA c) Give the most probable reaction for the electrochemical cell and the value for the cell s E o 2 Ag + S H + Ag 2 S + H 2 E o cell = V d) Draw a fully labeled diagram of the cell. e) As this reaction proceeds, what will happen to the E o value? Voltage will decrease because concentration of reactants is decreasing over time. f) What would happen if HCl(aq) was added to the cathodic halfcell? If HCl was added, then [H+] would increase, shifting the half reaction to the products.

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