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Redox potentials
Electrochemistry
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Redox ReactionsOxidation
• loss of electrons
Reduction
• gain of electrons
oxidizing agent
• substance that cause oxidation by being reduced
reducing agent
• substance that cause oxidation by being reduced
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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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Voltaic Cells
• harnessed chemical reaction which produces an electric current
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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) + 2 e-
reduction half reaction
cathode Cu+2(aq) + 2 e- ---> Cu(s)
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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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Voltaic Cells
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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) + 2 e-
reversible reaction
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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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Some Standard Reduction Potentials Table 18-1, pg 837
Li+ + e- ---> Li -3.045 v
Zn+2 + 2 e- ---> Zn -0.763v
Fe+2 + 2 e- ---> Fe -0.44v
2 H+(aq) + 2 e- ---> H2(g) 0.00v
Cu+2 + 2 e- ---> Cu +0.337v
O2(g) + 4 H+(aq) + 4 e- ---> 2 H2O(l) +1.229v
F2 + 2e- ---> 2 F- +2.87v
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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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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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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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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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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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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+?
yes, no
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Gibbs Free Energyand Cell Potential
G = - nFE
where n => number of electrons changed
F => Faraday’s constant
E => cell potential
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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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Applications of Electrochemical Cells
Batteries
Primary Cells
"dry" cell & alkaline cell 1.5 v/cell
mercury cell 1.34 v/cell
fuel cell 1.23v/cell
Secondary Cells
lead-acid (automobile battery) 2 v/cell
NiCad 1.25 v/cell
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“Dry” Cell
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“Dry” Cell
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“Flash Light” Batteries
"Dry" CellZn(s) + 2 MnO2(s) + 2 NH4
+ ----->
Zn+2(aq) + 2 MnO(OH)(s) + 2 NH3
Alkaline CellZn(s) + 2 MnO2(s) ---> ZnO(s) + Mn2O3(s)
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“New” Super Iron Battery
Mfe(VI)O4 + 3/2 Zn 1/2 Fe(III)2O3 + 1/2 ZnO + MZnO2
(M = K2 or Ba)Environmentally friendlier than MnO2 containing batteries.
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Lead-Acid(Automobile Battery)
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Lead-Acid(Automobile Battery)
Pb(s) + PbO2(s) + 2 H2SO4 = 2 PbSO4(s) + 2 H2O
2 v/cell
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Nickel-Cadmium (Ni-Cad)
Cd(s) + 2 Ni(OH)3(s) = Cd(OH)2(s) + 2 Ni(OH)2(s)
NiCad 1.25 v/cell
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Automobile Oxygen Sensor
Air, constant [O ]2
porous Pt electrodes
migrating O ions2-
measured potential difference
exhaust gas, unknown [O ]2
ZrO / CaO2
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Automobile Oxygen Sensor
• see Oxygen Sensor Movie from Solid-State Resources CD-ROM
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pH = (Eglass electrode - constant)/0.0592
pH Meter
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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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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 = (1.10 - (0.0257/2)ln 10) V
Ecell = (1.10 - (0.0257/2)2.303) V
Ecell = (1.10 - 0.0296) V = 1.07 V
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Nernst Equation
+0.83 V
salt bridge
H in2
1 atm
voltmeter
Pt electrode
Pt electrodeNaOH
1 M
H in2
1 atm
e– e–
anode (–)
HCl 1 M
cathode (+)
pn+
+
+
+
–
–
–
–
[H+]acid side [H+]base side
E = Eo – RTnF ln Q =
– 2.3 RTF log
[H+]base side[H+]acid side
[h+]p-type side [h+]n-type side
E (in volts) = – 2.3 RT
F log [h+]n-type side [h+]p-type side