2. 1 ELECTROCHEMISTRY C H A P T E R 1 1

3. 2 ELECTROCHEMISTRY: RELATIONSHIP OF ELECTRICAL CHARGE OR ELECTRICITY TO CHEMICAL REACTIONS REDUCTION + OXIDATIONOX ELECTRONS FROM DONOR ORBITALS OF A REDUCING AGENT TRANSFER TO THE ACCEPTOR ORBITALS OF THE OXIDIZING AGENT! ENERGYENERGY DONOR ACCEPTOR SPONTANEOUSNON-SPONTANEOUS OR ELECTRON TRANSFER REACTIONS

4. 3 ENERGYENERGY DONOR ACCEPTOR SPONTANEOUSNON-SPONTANEOUS OXIDATION = LOSS OF ELECTRONS ELECTRONS MOVE FROM HIGHER TO LOWER ENERGYDONOR ORBITALS ARE FULL OR PARTIALLY FULL = INCREASE IN OXIDATION STATE SPECIES OXIDIZED = REDUCING AGENT OR REDUCTANT = ANODIC PROCESS REDUCTION= GAIN OF ELECTRONS= DECREASE IN OXIDATION STATE ELECTRONS MOVE FROM LOWER TO HIGHER ENERGYACCEPTOR ORBITALS ARE UNFILLEDSPECIES REDUCED = OXIDIZING AGENT OR OXIDANT = CATHODIC PROCESS GALVANIC CELL ELECTROLYTIC CELL

5. 4 REDUCTION:  GAIN OF ELECTRON(S) Fe 3+ + 1e - Fe 2+ + 2e - Fe REDUCE THE CHARGE  CATHODIC PROCESS  OXIDIZING AGENT OR OXIDANT OXIDATION:  LOSS OF ELECTRON(S) Sn Sn 2+ + 2e - Sn 4+ + 2e - INCREASE THE CHARGE  ANODIC PROCESS  REDUCING AGENT OR REDUCTANT

6. 5 ENERGYENERGY DONOR: HIGHER ENERGY GOOD REDUCING AGENT IS OXIDIZED ACCEPTOR LOWER ENERGY GOOD OXIDIZING AGENT IS REDUCED SPONTANEOUS OR  G <0 NON-SPONTANEOUS OR  G > 0  G ~ E ACCEPTOR - E DONOR  G < 0 IF E DONOR < E ACCEPTOR

7. 6 Cu Zn SALT BRIDGE 1 M CuSO 4 1 M ZnSO 4 CATHODE ANODE e-e- + - K 1+ Cl 1- Zn 2+ Zn Cu 2+ Cu 1.10 V Cu 2+ + 2e - Cu Zn 2+ + 2e - Zn Cu 2+ + Cu + Zn 2+ Zn

8. 7 Cu 2+ + 2e - Cu Zn 2+ + 2e - Zn Cu 2+ + Cu + Zn 2+ Zn  HALF REACTIONS OR HALF CELLS REDUCTION HALF REACTION OXIDATION HALF REACTION REDOX REACTION NOTE THAT IN A BALANCED REDOX REACTION: THERE IS AN ATOMIC/IONIC/MOLECULAR BALANCE THERE IS A CHARGE BALANCE!!!!! ANY REACTION IN WHICH THERE IS A CHANGE IN AN OXIDATION STATE INVOLVES A REDOX REACTION

9. 8 BALANCING REDOX REACTIONS Fe 2+ + Cr 2 O 7 2- Fe 3+ + Cr 3+ (ACIDIC SOL’N) Fe 2+ Fe 3+ Cr 2 O 7 2- 2Cr 3+ + 1e - + 6e - + 7H 2 0 14H + + KEY: NET TRANSFER OF ELECTRONS MUST = O X6 6 6 Cr 2 O 7 2- 2Cr 3+ + 7H 2 0 + 6Fe 3+ 14H + + Fe 2+ + 6 TOTAL ELCETRON TRANSFER = 6

10. 9 As (s) + ClO 3 1- (aq)  H 3 AsO 3 (aq) + HClO (aq) ClO 3 1-  HClO 0 +3 + 3e - +5 +1 4e - + + 2H 2 O5H 1+ + X3 X4 As  H 3 AsO 3 + 3H 1+ 3H 2 O + 4As + 3ClO 3 1-  4 H 3 AsO 3 + 3 HClO 12H 2 O + + 12H 1+ 15H 1+ + + 6H 2 O 4As + 3ClO 3 1-  4 H 3 AsO 3 + 3 HClO 6H 2 O + 3H 1+ + As  H 3 AsO 3 + 3H 1+ 3H 2 O ++ 3e - ClO 3 1-  HClO 4e - ++ 2H 2 O5H 1+ + OXIDATION REDUCTION ANODE CATHODE

11. 10  DEFINE THE FARADAY, F: CHARGE PER MOLE e - OR COULOMBS/MOL F = 9.65 x 10 4 COULOMBS/MOL e - FOR n MOLES: - n F  VOLTAGE OR POTENTIAL DIFFERENCE, E ENERGY/CHARGE OR JOULE/CHARGE GALVANIC CELL PROVIDES ELECTRICITY OR ELECTRONS TO AN OUTSIDE CIRCUIT!  WORK DONE BY ELECTRONS = - ( - n F E )  WORK DONE BY ELECTRONS =  G  G = - n F E  AND  G o = - n F E o

12. 11 CELL POTENTIALS (VOLTAGES) GALVANIC CELL: ELECTROCHEMICAL CELLS IN WHICH ELECTRONS TRANSFER SPONTANEOUSLY BATTERY  G < 0 Mg 2+ +2e - Mg Ag 1+ +e - Ag E o = -2.356 V vs NHE E o = +0.800 V vs NHE OXIDATION REDUCTION ANODE CATHODE E o cell = E o cathode - E o anode E o cell = 0.800 - (-2.356)3.156 V 2Ag 1+ (aq) + Mg  2Ag + Mg 2+ (aq) Ag 1+ (aq) + Mg  Ag + Mg 2+ (aq) GoGo = - n F E o SPONTANEOUS DIRECTION: E o cathode > E o anode GOOD REDUCING AGENT: LOW E o -2(96500)(3.156) - 609108 OR - 609 kJ/mol

13. 12 14H + +Cr 2 O 7 2-  2Cr 3+ + 7H 2 0 + 6Fe 3+ Fe 2+ + 6 6Fe 2+ 6Fe 3+ + 6e - 14H 1+ + Cr 2 O 7 2- + 6e - 2Cr 3+ + 7H 2 O OXIDATION REDUCTION ANODE CATHODE E o cell = E o cathode - E o anode LOOK UP: Fe 3+ + e - Fe 2+ E o = +0.77 V 14H 1+ + Cr 2 O 7 2- 2Cr 3+ + 7H 2 O E o = +1.36 V = +1.33 - (+0.77)= 0.56 V GoGo = - n F E o  G o = -324240 J/mol OR -324 kJ/mol =-6 (96500)(0.56)

14. 13 GoGo = - n F E o = -212300 J/mol OR -212 kJ/mol E o cell = E o cathode - E o anode LOOK UP: Cu 2+ + 2eCu Zn 2+ + 2e Zn E o = +0.34 V E o = -0.76 V REDUCTION (CATHODE) OXIDATION (ANODE) = +0.34 - (-0.76)= 1.10 V = - 2(96500)(1.10) Cu 2+ + Cu + Zn 2+ Zn

15. 14 E o cell = E o cathode - E o anode Pb 2+ + 2e Pb Fe 2+ + 2e Fe E o = -0.13 V E o = -0.44 V Fe 2+ + 2e Fe Al 3+ + 3e Al E o = -1.66 V E o = -0.44 V E o cell = E o cathode - E o anode = -0.13 - (-0.44)= 0.31 V Pb 2+ + Fe Pb + Fe 2+ = -0.44 - (-1.66) 3Fe 2+ + 2Al 3Fe + 2Al 3+ = 1.22 V

16. 15 CORROSION: GALVANIC CELL REEACTION METAL REDUCING AGENT REACTION WITH H 2 O AND O 2 OR S O 2 + 4H 1+ + 4e -  2H 2 O 2Fe  2Fe 2+ + 4e - CATHODE ANODE E o cell = +1.67 V 2 nd OXIDATION OF Fe 2+  Fe 3+ Fe 2 O 3 OTHER FORM OF “BAD” CORROSION: SILVER TARNISH Fe 2 O 3

17. 16 TO PREVENT CORROSION: REMOVE H 2 O, O 2 OR BOTH! SACRIFICIAL ANODE GALVANIZATION PAINT PLATING CONVERSION COATINGS NATURAL PASSIVATION Al 2 O 3, SnO 2, STAINLESS STEEL, MgO, PATINA“GOOD” CORROSION ELECTROLYTIC CELLS

18. 17 CHEMICAL ENERGY ELECTRICAL ENERGY GALVANIC CELL ELECTROLYTIC CELL