Download Voltaic Cells

Redox potentials
Electrochemistry
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Redox Reactions
Oxidation
• 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 ereduction 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 ereversible 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 halfreaction in which reduction is occurring
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Some Standard Reduction Potentials
Table 18-1, pg 837
Li+ + e- ---> Li
Zn+2 + 2 e- ---> Zn
Fe+2 + 2 e- ---> Fe
2 H+(aq) + 2 e- ---> H2(g)
Cu+2 + 2 e- ---> Cu
O2(g) + 4 H+(aq) + 4 e- ---> 2 H2O(l)
F2 + 2e- ---> 2 FICS
-3.045 v
-0.763v
-0.44v
0.00v
+0.337v
+1.229v
+2.87v
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 halfcells 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 Energy
and Cell Potential
DG = - 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" Cell
Zn(s) + 2 MnO2(s) + 2 NH4+ ----->
Zn+2(aq) + 2 MnO(OH)(s) + 2 NH3
Alkaline Cell
Zn(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
porous Pt electrodes
ZrO2 / CaO
exhaust gas,
unknown [O 2 ]
migrating
O2- ions
Air,
constant [O2 ]
measured potential difference
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Automobile Oxygen Sensor
• see Oxygen Sensor Movie from Solid-State
Resources CD-ROM
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pH Meter
pH = (Eglass electrode - constant)/0.0592
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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
[H+ ]acid side [H+ ]base side
E = Eo –
[H+ ]base side
RT
– 2.3 RT
ln
=
log +
nF Q
F
[H ]acid side
[h+ ]p-type side  [h+ ]n-type side
E (in volts) =
[h+ ]n-type side
– 2.3 RT
log +
F
[h ]p-type side
voltmeter
e–
+0.83 V
e–
salt bridge
H2 in
H2 in
1 atm
1 atm
+ –
n
+ –
+ –
+ –
Pt
electrode
NaOH
1M
anode (–)
HCl
1M
cathode (+)
Pt
electrode
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p