Batteries and Electrolysis

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Batteries and Electrolysis

• CHEM 1120
• April 13, 2005

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Dry Cell and Alkaline Battery
Gallery (p. 924)
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Chemistry of BatteriesI
Dry Cells
Anode (oxidation) Zn(s) Zn2(aq) 2
e- Cathode (reduction) 2 MnO2 (s) 2 NH4(aq)
2 e- Mn2O3 (aq) 2 NH3 (aq)
H2O(l) ammonia is tied up with Zn2
Zn2(aq) 2 NH3 (aq) 2 Cl-(aq)
Zn(NH3)2Cl2 (s) Overall (cell) reaction 2 MnO2
(s) 2NH4Cl(aq) Zn(s) Zn(NH3)2Cl2 (s)
H2O(l) Mn2O3 (s)
Ecell 1.5 V
Alkaline Battery
Anode (oxidation) Zn(s) 2 OH-(aq)
ZnO(s) H2O(l) 2 e - Cathode
(reduction) MnO2 (s) 2 H2O(l) 2 e -
Mn(OH)2 (s) 2 OH -(aq) Overall (cell)
reaction Zn(s) MnO2 (s) H2O(l)
ZnO(s) Mn(OH)2 (s) Ecell 1.5 V
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Gallery (p. 925)
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Chemistry of BatteriesII
Mercury and Silver (Button) Batteries
Anode (oxidation) Zn(s) 2 OH-(aq)
ZnO(s) H2O(l) Cathode (reduction)
(mercury) HgO(s) H2O(l) 2 e- Hg(l)
2 OH-(aq) Cathode (reduction) (silver) Ag2O(s)
H2O(l) 2 e- 2 Ag(s) 2
OH-(aq) Overall (cell) reaction (mercury) Zn(s)
HgO(s) ZnO(s) Hg(l) Ecell
1.3 V Overall (cell) reaction (silver) Zn(s)
Ag2O(s) ZnO(s) 2 Ag(s)
Ecell 1.6 V
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Gallery (p. 925)
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Chemistry of BatteriesIII
Secondary (Rechargeable) Batteries
Lead-Acid Batteries Anode (oxidation)
Pb(s) SO42-(aq) PbSO4 (s) 2
e- Cathode (reduction) PbO2 (s) 4
H(aq) SO42-(aq) 2 e- PbSO4
(s) 2 H2O(l) Overall (cell) reaction
(discharge) PbO2 (s) Pb(s) 2 H2SO4 (aq)
2 PbSO4 (s) 2 H2O(l) Ecell 2 V
Overall cell) reaction (recharge) 2 PbSO4
(s) 2 H2O(l) PbO2 (s) Pb(s)
2 H2SO4 (aq)
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Lithium Solid-State Battery
Nickel-Cadmium (Nicad) Battery
Gallery (p. 926)
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Chemistry of BatteriesIV
Secondary (Rechargeable) Batteries (Continued)
Nickel-Cadmium (Nicad) Battery
Anode (oxidation) Cd(s) 2 OH-(aq)
Cd(OH)2 (s) 2 e- Cathode (reduction) 2
NiO(OH)(s) 2 H2O(l) 2 e- 2
Ni(OH)2 (s) 2 OH-(aq) Overall (cell)
reaction Cd(s) 2 NiO(OH)(s) 2 H2O(l)
2 Ni(OH)2 (s) Cd(OH)2 (s)
Ecell 1.4 V
Lithium Solid-State Battery
Anode (oxidation) Li(s) Li
(in solid electrolyte) e- Cathode (reduction)
MnO2 (s) Li e- LiMnO2
(s) Overall (cell) reaction Li(s) MnO2 (s)
LiMnO2 (s) Ecell 3 V
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The Tin-Copper Reaction as the Basis of a Voltaic
and an Electrolytic Cell
Fig. 21.17
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Construction and Operation of an Electrolytic Cell
Lets use the tin-copper voltaic cell shown in
Fig 21.17A
Sn(s) Sn2(aq) 2
e- anode oxidation Cu2(aq) 2 e-
Cu(s)
cathode reduction
Sn(s) Cu2(aq) Sn2(aq) Cu(s)
Eocell 0.48 V and Go -93 kJ
The spontaneous reaction of Sn metal through
oxidation to form Sn2ions and the reduction of
Cu2 ions to form copper metal will produce a
cell voltage of 0.48 volt. Therefore, the
reverse reaction is nonspontaneous and never
happens on its own. We can make the reverse
reaction occur by supplying power from an
external source with an electric potential
greater than Eocell. We convert the voltaic cell
into an electrolytic cell and reverse the
electrodes.
Cu(s) Cu2(aq)
2e- anode oxidation Sn2(aq) 2 e-
Sn(s)
cathode reduction
Sn2(aq) Cu(s) Cu2(aq) Sn(s)
Eocell -0.48 V and Go 93 kJ
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The Processes Occurring During the Discharge and
Recharge of a Lead-Acid Battery
Fig. 21.18
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Comparison of Voltaic and Electrolytic Cells
Electrode
Cell Type G Ecell
Name Process Sign
Voltaic lt 0 gt 0
Anode Oxidation - Voltaic
lt 0 gt 0
Cathode Reduction
Electrolytic gt 0 lt 0
Anode Oxidation
Electrolytic gt 0 lt 0
Cathode Reduction -
Table 21.4 (p. 932)
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The Electrolysis of Water
Fig. 21.19
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Uses of Electrolysis

• electrolysis of water to produce hydrogen and
  oxygen gases
• 2 H2O (l) ? 2 H2 (g) O2 (g)
•   refinement of metals from natural ores
• refinement of metals from impure samples
• electroplating to decorate and protect the
  surface of metals or other objects that have been
  coated with graphite to increase their electrical
  conductance