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 CuS Cu SO42-
b NO 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 NO NO + 2H2O

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

8. I. Balancing Redox Reactions (continued)
STEP 5. Balance Charge by inserting Electrons, where necessary
a5. CuS + 4H2O Cu2+ + SO H+ + 8e-
b5. NO H+ + 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 3Cu SO H e-
b6.Multiply by 8x
8NO H+ + 24e NO + 16H e-

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

11. I. Balancing Redox Reactions (continued)
So, the final, balanced reaction is:
3CuS(s) + 8 NO3-(aq) + 8H+ (aq) Cu2+(aq) + 3 SO42-(aq) + 8NO(g) H2 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 C for pure Al2O3)
Graphite Anodes (+)
Steel case
C lining
(Cathode)
(-) Al
Al2O3 in molten Na3AlF6 Al
Molten Al

17. Hall Cell for Al Manufacture

18. 2 Al2O3 (sln) + 3C (s) 4 Al (l) + 3CO2 (g)
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.