Electrochemistry

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Electrochemistry

• And REDOX Reactions

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Voltaic Cell
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Copper-Zinc Reaction
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A voltaic cell based on the zinc-copper reaction
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Notation for a Voltaic Cell
components of anode compartment (oxidation
half-cell)
components of cathode compartment (reduction
half-cell)
phase of lower oxidation state
phase of lower oxidation state
phase of higher oxidation state
phase of higher oxidation state
phase boundary between half-cells
Examples
Zn(s) Zn2(aq) Cu2(aq) Cu (s)
graphite I-(aq) I2(s) H(aq), MnO4-(aq)
Mn2(aq) graphite
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Determining an unknown E0half-cell with the
standard reference (hydrogen) electrode.
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Writing Spontaneous Redox Reactions

• By convention, electrode potentials are written
  as reductions.

• When pairing two half-cells, you must reverse one
  reduction half-cell to produce an oxidation
  half-cell. Reverse the sign of the potential.

• The reduction half-cell potential and the
  oxidation half-cell potential are added to obtain
  the E0cell.

• When writing a spontaneous redox reaction, the
  left side (reactants) must contain the stronger
  oxidizing and reducing agents.

stronger reducing agent
weaker oxidizing agent
stronger oxidizing agent
weaker reducing agent
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Selected Standard Electrode Potentials
(298K)
Half-Reaction
E0(V)
2.87
1.36
1.23
0.96
0.80
0.77
0.40
0.34
0.00
2H(aq) 2e- H2(g)
-0.23
-0.44
-0.83
-2.71
-3.05
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Calculating Eocell

• The EMF for the cell is calculated as follows
• Eocell Ered - Eox
• More explicitly, the voltage of the cell is
  equal to the half cell voltage of the reduction
  reaction minus the half-cell voltage of the
  oxidation reaction

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Examples

• Determine the EMF for each of the following
  cells, using potentials from the appendix in the
  back of your text
• Fe3 Fe2
• Zn2 Zn(s)
• PbO2(s) HSO4- 3H PbSO4(s) 2 H2O
• PbSO4(s) H Pb(s) HSO4-
• MnO4- 2H2O MnO2 4 OH-
• Cl2 2Cl-

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2
3
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Summary of Relationship Among Variables
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Dry Cell

• ZincCarbon Dry Cell Alkaline
• Anode Zn(s) 2OH-(aq) ? Zn(OH)2(s) 2e
• Cathode 2MnO2(s) H2O(l) 2e- ? Mn2O3(s)
  2OH- (aq)
• This cell performs better under current drain
  and in cold weather. It isnt truly dry but
  rather uses an aqueous paste.

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Dry Cell Design
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Lead Acid Storage Battery

• Lead Storage Cell
• The electrodes are lead alloy grids one is
  packed with a spongy lead to form the anode, and
  the other is packed with lead dioxide to form the
  cathode. Both electrodes are in an aqueous
  solution of H2SO4.
• Anode Pb(s) HSO4-(aq) ? PbSO4(s) H(aq)
  2e-
• Cathode PbO2(s) 3H(aq) HSO4-(aq) 2e- ?
  PbSO4(s) 2H2O(l)
• Unlike dry cells, after discharge, lead storage
  cells can be recharged.

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Lead Acid Battery Design
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Recharging a Lead Acid Battery
Each of the six cells in the lead storage cell
generates 2 V, yielding 12 V. During discharge,
white PbSO4(s) coats each electrode. To recharge
the cell, an external current is used, reversing
the previous reactions. Some water decomposes
into hydrogen and oxygen gas, so more water may
need to be added. Newer batteries use electrodes
with calcium in the lead, which resists
decomposition by water. These versions are
maintenance free.
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Ni-Cad Batteries
NickelCadmium Cell Anode Cd(s) 2OH-(aq) ?
Cd(OH)2(s) 2e- Cathode NiOOH(s) H2O(l) e-?
Ni(OH)2(s) OH-(aq)

• These cells are used in calculators, portable
  power tools, shavers, and toothbrushes. During
  recharge, the reactions are reversed, which can
  be done many times.

• When cadmium is replaced with a metal hydride
  (MH), nickel metal hydride and lithium hydride
  cells result. They are less toxic.

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Fuel Cells

• Fuel Cell
• Fuel cells require a continuous supply of
  reactants (fuel).
• Anode H2(g) ? 2H(aq) 2e-
• Cathode O2(g) 4H(aq) 4e- ? 2H2O(l)
• Fuel cells were originally used in space
  applications, but are now being explored for more
  uses.

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Fuel Cell Design
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Corrosion (rusting)
Anode Fe(s) ? Fe2(aq) 2e- Cathode O2(g)
2H2O(l) 4e- ? 4OH-(aq)
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Corrosion Control
Corrosion Control Cathodic Protection Voltaic
cells can be used to control corrosion of
underground pipelines and tanks. Rusting occurs
when water comes in contact with iron. The edge
of the water drop, when exposed to air, becomes
one pole of a voltaic cell where oxygen is
reduced to hydroxide. When the buried metal is
connected to a more active metal such as
magnesium, the magnesium becomes the anode and
the iron becomes the cathode. The iron is,
therefore, protected from oxidation. This
phenomenon is called cathodic protection.
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Corrosion Control
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Corrosion Control
Another method of corrosion control for iron is a
zinc coating. Placed on the iron in a hot-dip
procedure, the zinc, which is more active than
iron, acts as a sacrificial barrier since it gets
oxidized before the iron. The zinc oxide coating
that forms eventually reacts with air to become a
very tough, durable coating composed of zinc
carbonate. Iron treated in this fashion is said
to be galvanized.
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Electrolysis

• Electrolytic Cell
• An electrolytic cell is an electrochemical cell
  in which an electric current drives an otherwise
  nonspontaneous reaction.
• The process of producing a chemical change in an
  electrolytic cell is called electrolysis. Many
  important substances are produced commercially by
  electrolysisfor example, aluminum and chlorine.

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Electrolysis

• Downs Cell
• A Downs cell is an electrolytic cell used to
  obtain sodium metal by electrolysis of sodium
  chloride. The products must be kept separated or
  they would react.
• Anode Cl-(l) ? ½Cl2(g) e-
• Cathode Na(l) e- ? Na(l)

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Electrolysis
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Electrolysis