Redox Reactions & Electrochemical Cells

1. Redox Reactions & Electrochemical Cells I. Balancing Redox Reactions

2. I.Balancing Redox Reactions • STEP 1. Split Reaction into 2 Half-Reactions • STEP 2. Balance Elements Other than H & O • STEP 3. Balance O by Inserting H2O into eqns. as necessary • STEP 4. Balance H with H+ or H2O (see 4a, 4b) • STEP 5. Balance Charge by Inserting Electrons as needed • STEP 6. Multiply Each 1/2 Reaction by Factor needed to make no. of Electrons in each 1/2 Reaction Equal • STEP 7. Add Eqns. & Cancel Out Duplicate terms, where possible

3. I.Balancing Redox Reactions (continued) • STEP 4a. In ACID: Balance H by Inserting H+, as needed • STEP 4b. In BASE: Balance H by (i) inserting 1 H2O for each missing H & (ii) inserting same no. of OH- on OTHER SIDE OF REACTION as H2Os added in (i)

4. I.Balancing Redox Reactions (continued) • Example • Complete and Balance Following Reaction: • CuS (s) + NO3 -(aq) Cu2+(aq) + SO42-(aq) • + NO (g) • STEP1. Split into 2 Half-Reactions • a.1 CuS Cu2+ + SO42- • b.1 NO3 - NO

5. I.Balancing Redox Reactions (continued) • STEP 2. Balance Elements Other than H & O • Already O.K. !

6. I.Balancing Redox Reactions (continued) • STEP 3. Balance O by inserting H2O into equations as necessary • a.3 CuS + 4H2O Cu2+ + SO42- • b.3 NO3- NO + 2H2O

7. I.Balancing Redox Reactions (continued) • STEP 4. ACIDIC, so Balance H by inserting H+ as needed • a4. CuS + 4H2O Cu2+ + SO42- + 8H+ • b4. NO3- + 4H+ NO + 2H2O

8. I.Balancing Redox Reactions (continued) • STEP 5. Balance Charge by inserting Electrons, where necessary • a5. CuS + 4H2O Cu2+ + SO42- + 8H+ + 8e- • b5. NO3- + 4H+ + 3e- NO + 2H2O

9. I.Balancing Redox Reactions (continued) • STEP 6. Multiply each Eqn. by factor to make No. of Electrons in Each 1/2 Reaction the Same • a6. Multiply by 3x • 3CuS + 12H2O 3Cu2+ + 3SO42- + 24H+ + 24e- • b6.Multiply by 8x • 8NO3- + 32H+ + 24e- 8NO + 16H+ + 24e-

10. I.Balancing Redox Reactions (continued) • STEP 7. Add Eqns. and Cancel Out Duplicated Terms • (a7 + b7) 8H+ • 3CuS + 12H2O + 8NO3- + 32H+ + 24 e- • 3Cu2+ + 3SO42- + 24H+ + 8NO +16 H2O +24e- 4H2O

11. I.Balancing Redox Reactions (continued) So, the final, balanced reaction is: 3CuS(s) + 8 NO3-(aq)+ 8H+(aq) 3Cu2+(aq) + 3 SO42-(aq) + 8NO(g) + + 4H2 O(l)

12. Checking mass balance and charge balance in Equation • L.H.S • 3 x Cu • 3 x S • 8 x N • 24 x O • 8 x H • (8 x 1-) + (8 x H+) = 0 • R.H.S. • 3 x Cu • 3 x S • 8 x N • 24 x O • 8 x H • (3 x 2+ )+(3 x 2- ) = 0

13. Redox Reactions in Electrochemistry • Two Types of Electrochemical Cells: • 1. Galvanic • 2. Electrolytic • Galvanic Cell - Converts a Chemical Potential Energy into an Electrical Potential to Perform Work • Electrolytic Cell- Uses Electrical Energy to Force a Chemical Reaction to happen that would not otherwise occur

14. Anode and Cathode in Electrochemistry • ANODE - Where OXIDATION takes place • (-e-) • CATHODE - Where REDUCTION takes place (+e-)

15. Electrochemistry and the Metals Industry • Many Electrochemical Processes are used Commercially for Production of Pure Metals: • e.g. Al Manufacture (by electrolysis of Al2O3) • Mg Manufacture (by electrolysis of MgCl2) • Na Manufacture (by electrolysis of NaCl)

16. Electrolylitic Production of Al using the HALL CELL (major plant in ALCOA, TN • Al2O3 dissolved in molten cryolite (Na3AlF6) at 950 0C (vs. 2050 0C for pure Al2O3) Graphite Anodes (+) C lining (Cathode) (-) Steel case Al Al2O3 in molten Na3AlF6 Al Molten Al

17. Hall Cell for Al Manufacture

18. Hall Cell Process Reaction: 2 Al2O3(sln) + 3C (s) 4 Al (l) + 3CO2 (g) Location of Hall cell plant in E. Tennessee through availability of inexpensive Hydroelectric power. Process uses 50,000 – 100,000 A.