1. Redox potentials
Electrochemistry
ICS

2. Redox Reactions Oxidation loss of electrons Reduction
gain of electrons
oxidizing agent
substance that cause oxidation by being reduced
reducing agent
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3. Electrochemistry
In the broadest sense, electrochemistry is the study of chemical reactions that produce electrical effects and of the chemical phenomena that are caused by the action of currents or voltages.
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4. ICS

5. Voltaic Cells
harnessed chemical reaction which produces an electric current
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6. Voltaic Cells Cells and Cell Reactions Daniel's Cell
Zn(s) + Cu+2(aq) ---> Zn+2(aq) + Cu(s)
oxidation half reaction
anode Zn(s) ---> Zn+2(aq) e-
reduction half reaction
cathode Cu+2(aq) e- ---> Cu(s)
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7. Voltaic Cells
copper electrode dipped into a solution of copper(II) sulfate
zinc electrode dipped into a solution of zinc sulfate
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8. Voltaic Cells
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9. Hydrogen Electrode
consists of a platinum electrode covered with a fine powder of platinum around which H2(g) is bubbled. Its potential is defined as zero volts.
Hydrogen Half-Cell
H2(g) = 2 H+(aq) e-
reversible reaction
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10. ICS

11. ICS

12. Standard Reduction Potentials
the potential under standard conditions (25oC with all ions at 1 M concentrations and all gases at 1 atm pressure) of a half-reaction in which reduction is occurring
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13. Some Standard Reduction Potentials Table 18-1, pg 837
Li+ + e- ---> Li v
Zn e- ---> Zn v
Fe e- ---> Fe v
2 H+(aq) e- ---> H2(g) v
Cu e- ---> Cu v
O2(g) H+(aq) e- ---> 2 H2O(l) v
F e- ---> 2 F v
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14. If the reduction of mercury (I) in a voltaic cell is desired, the half reaction is:
Which of the following reactions could be used as the anode (oxidation)?
A, B
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15. Cell Potential
the potential difference, in volts, between the electrodes of an electrochemical cell
Direction of Oxidation-Reduction Reactions
positive value indicates a spontaneous reaction
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16. Standard Cell Potential
the potential difference, in volts, between the electrodes of an electrochemical cell when the all concentrations of all solutes is 1 molar, all the partial pressures of any gases are 1 atm, and the temperature at 25oC
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17. Zn(s)/ZnSO4(aq)//CuSO4(aq)/Cu(s)
Cell Diagram
the shorthand representation of an electrochemical cell showing the two half-cells connected by a salt bridge or porous barrier, such as:
Zn(s)/ZnSO4(aq)//CuSO4(aq)/Cu(s)
anode cathode
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18. Metal Displacement Reactions
solid of more reactive metals will displace ions of a less reactive metal from solution
relative reactivity based on potentials of half reactions
metals with very different potentials react most vigorously
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19. Ag+ + e- --->Ag E°= 0.80 V Cu2+ + 2e- ---> Cu E°= 0.34 V
Will Ag react with Cu2+?
yes, no
Will Cu react with Ag+?
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20. Gibbs Free Energy and Cell Potential
DG = - nFE
where n => number of electrons changed
F => Faraday’s constant
E => cell potential
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21. Applications of Electrochemical Cells
Batteries
device that converts chemical energy into electricity
Primary Cells
non-reversible electrochemical cell
non-rechargeable cell
Secondary Cells
reversible electrochemical cell
rechargeable cell
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22. Applications of Electrochemical Cells
Batteries
Primary Cells
"dry" cell & alkaline cell 1.5 v/cell
mercury cell v/cell
fuel cell 1.23v/cell
Secondary Cells
lead-acid (automobile battery) 2 v/cell
NiCad v/cell
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23. “Dry” Cell
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24. “Dry” Cell
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25. “Flash Light” Batteries
"Dry" Cell
Zn(s) MnO2(s) NH >
Zn+2(aq) MnO(OH)(s) NH3
Alkaline Cell
Zn(s) MnO2(s) ---> ZnO(s) + Mn2O3(s)
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26. “New” Super Iron Battery
Mfe(VI)O4 + 3/2 Zn /2 Fe(III)2O3 + 1/2 ZnO + MZnO2
(M = K2 or Ba)
Environmentally friendlier than MnO2 containing batteries.
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27. ICS

28. Lead-Acid (Automobile Battery)
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29. Lead-Acid (Automobile Battery)
Pb(s) + PbO2(s) H2SO4 = 2 PbSO4(s) H2O
2 v/cell
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30. Nickel-Cadmium (Ni-Cad)
Cd(s) Ni(OH)3(s) = Cd(OH)2(s) Ni(OH)2(s)
NiCad v/cell
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31. ICS

32. Automobile Oxygen Sensor
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33. Automobile Oxygen Sensor
see Oxygen Sensor Movie from Solid-State Resources CD-ROM
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34. pH Meter
pH = (Eglass electrode - constant)/0.0592
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35. Effect of Concentration on Cell Voltage: The Nernst Equation
Ecell = Eocell - (RT/nF)ln Q
Ecell = Eocell - (0.0592/n)log Q
where Q => reaction quotient
Q = [products]/[reactants]
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36. EXAMPLE: What is the cell potential for the Daniel's cell when the [Zn+2] = 10 [Cu+2] ? Q = ([Zn+2]/[Cu+2] = (10 [Cu+2])/[Cu+2] = 10 Eo = (0.34 V)Cu couple + (-(-0.76 V)Zn couple n = 2, 2 electron change Ecell = Eocell - (0.0257/n)ln Q
thus Ecell = ( (0.0257/2)ln 10) V
Ecell = ( (0.0257/2)2.303) V
Ecell = ( ) V = V
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37. Nernst Equation Q [H ] ® E = – RT nF ln – 2.3 F log [h ] ® E+ ]
acid side
base side E = o – RT nF ln Q
– 2.3  F
log
base side
acid side [h + ]
p-type side
n-type side E
(in volts) =
– 2.3  RT F
log
n-type side
p-type side
© 1993 American Chemical Society
In an electrochemical concentration cell, the cell potential depends on the concentrations of ions (in this case, H+ ion) in the two half-cells.
In semiconductors the voltage obtainable with a p-n junction depends on the relative concentration of electrons or holes on the two sides of the junction.
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