Chapter 23 Electrochemistry

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Chapter 23Electrochemistry
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Sections 23.1-23.2Electrochemical Cells

• OBJECTIVES
• Describe how RedOx rxns produce useful
  electricity
• Explain the structure and function of Voltaic
  (Galvanic) Cells i.e. batteries

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The Nature of Voltaic (Galvanic) Cells

• You have already seen that when a strip of metal
  is placed in a solution containing a less active
  metal, a single replacement rxn will occur
• This is a classic example of a RedOx rxn. Whether
  the process is spontaneous or not can easily be
  predicted by using Table J in your Reference
  Tables for Physical Setting / CHEMISTRY

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Voltaic/Galvanic Cells

• If the half-rxns which define a RedOx rxn are
  allowed to occur in separate beakers, the
  electrons can be made to flow through an external
  wire and used to perform work
• The problem is that as the RedOx rxn tries to
  proceed there is an imbalance of ions in the
  beakers. Nature will not allow this, so we must
  supply ions from an external source to keep each
  solution electrically neutral.
• The external source of ions is called a Salt
  Bridge. It is composed of a salt- saturated gel
  contained by a glass U-tube.

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The Daniell Cell
The Daniell Cell was the first wet cell
battery. It is composed of a copper electrode in
a copper (II) sulfate solution, a zinc electrode
in a zinc sulfate solution, a salt bridge
(usually containing sodium chloride), an external
wire, and a voltmeter. The electrons
spontaneously flow from the Zn electrode to the
Cu electrode. How the cell works is described in
the next slide.
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Daniell Cell Half-Rxns

• Oxidation Half-Rxn
• Zn(s) ? Zn2(aq) 2e-1 ANODE (-)
• Reduction Half-Rxn
• Cu2(aq) 2e-1 ? Cu(s) CATHODE ()
• An Ox, Red Cat

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So What Else Happens in the Daniell Cell?

• The electrons flow from the Zn electrode to the
  Cu electrode
• The two RedOx rxns happen simultaneously
• The cation in the salt bridge moves toward the
  cathode (Cu Electrode)
• The anion in the salt bridge moves toward the
  anode (Zn Electrode)
• The Zn electrode gets lighter in mass
• The Cu electrode gets heavier in mass
• The rxn stops (the battery is dead) when the salt
  in the salt bridge runs out, or the Zn electrode
  is used up, or the Cu2 ions run out

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Cell Voltages

• Standard cell conditions are defined as
• 1 Molar Solute, 25? C, 1 Atm
• A Standard Hydrogen Electrode under standard
  conditions has a back voltage applied so the
  observed cell voltage appears to be zero (see
  diagrams in the next slides)
• If a stated voltage (E) has a zero superscript
  (E?), the experiment was done under standard
  conditions

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(a) The Standard Hydrogen Cell and (b) Close up
of the Hydrogen Standard Electrode.
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Cell Voltages

• You text book has a table of Reduction Potentials
  (voltages) on pg. 688
• Find the two half reactions for your cell in
  this case we have
• Cu2(aq) 2e-1 ? Cu(s) E? 0.34 V
• Zn2 (aq) 2e-1 ? Zn (s) E? -0.76 V
• Since you must have both a Red and an Ox rxn,
  turn the half rxn with the smaller voltage around
  and change the sign of the voltage (next slide)

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Cell Voltages (continued)

• In the present case we have
• Cu2(aq) 2e-1 ? Cu(s) E? 0.34 V
• Zn(s) ? Zn2 (aq) 2e-1 E? 0.76
  V
• If the number of electrons on each side is the
  same simply add the half rxns together and
  simplify the voltages are also added together in
  a similar fashion
• Cu2(aq) Zn(s) ? Zn2 (aq) Cu(s) Overall
  Rxn
• 0.34 (0.76) 1.10 Volts E?cell

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The Daniell Cell under Standard
Conditions(notice the cell voltage!)
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Another Example (pg. 1 of 5)

• PROBLEM Suppose someone gave you Al(s), Zn(s),
  Al(NO3)3(aq), Zn(NO3)2(aq), and NaC2H3O2 as a
  paste construct a Voltaic cell and label or
  explain all components
• SOLUTION First, identify the electrodes (usually
  solids), then immerse them in their appropriate
  electrolytes, and let the paste be the salt
  bridge so in this case we have
• Al(s) in Al(NO3)3(aq)
• Zn(s), in Zn(NO3)2(aq)
• NaC2H3O2 as a paste is in the salt bridge

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Another Example (pg. 2 of 5)

• Now write the Reduction Half Rxns using pg. 688
• Al3(aq) 3e-1 ? Al(s) E?
  -1.66 V
• Zn2(aq) 2e-1 ? Zn(s) E?
  -0.76 V
• You must have the same number of electrons on
  both sides of the arrow, so multiple the first
  rxn by 2 and the second rxn by 3. The voltages,
  however, are not changed
• 2Al3(aq) 6e-1 ? 2Al(s) E?
  -1.66 V
• 3Zn2(aq) 6e-1 ? 3Zn(s) E?
  -0.76 V

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Another Example (pg. 3 of 5)

• Turn the smaller voltage rxn around, change sign,
  and add
• 2Al(s) ? 2Al3(aq) 6e-1 (Ox) E?
  1.66 V
• 3Zn2(aq) 6e-1 ? 3Zn(s) (Red) E? -0.76
  V
• 2Al(s) 3Zn2(aq) ? 3Zn(s) 2Al3(aq)
• Over All Net Ionic Rxn (balanced by charge
  mass!)
• E?cell 0.90 V (1.66 V) (-0.76 V)
• Note that all cell voltages must be positive!

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Another Example (pg. 4 of 5)

• Thus, we know
• Electrons flow from Al(s) ? Zn(s)
• The cell voltage is 0.90 V
• Reduction occurs at the Zn electrode, and it is
  the cathode ()
• Oxidation occurs at the Al electrode, and it is
  the anode (-)
• Na1(aq) from the salt bridge flows to the
  cathode (Zn)
• C2H3O2(aq)-1 from the salt bridge flows to the
  anode (Al)
• The Zn electrode gets heavier while the Al
  electrode gets lighter

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Another Example (pg. 5 of 5)
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Other Batteries Lead Storage Cell
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Other Batteries Standard Dry Cell
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Other Batteries Alkaline Battery
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Section 23.3Electrolytic Cells More

• OBJECTIVES
• Describe how RedOx rxns can be used to
  electroplate
• Discuss some other practical applications of
  electrochemistry

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Electrolytic Cell Basics

• Electrolytic Cells require an external D.C. power
  supply
• The power supply forces a RedOx rxn to take place
  backwards (rxn normally has a positive DG? or a
  negative E?)
• Usually, there is only one beaker containing both
  the electrolyte and electrode
• AnOx RedCat still works .. but the cathode is
  now (-) and the anode is now (). This is
  backwards compared to a Voltaic Cell!

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Electroplating of Silver Metal
The oxidation of sliver metal has an E? value of
-0.80 V, so it is not a spontaneous reaction. The
external power source supplies this voltage to
drive the rxn as seen in the adjacent diagram.
The silver electrode is oxidized (loses weight)
and the spoon is plated by the reduction of
silver ions.
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The Formation of Rust
Water acts as the medium for the RedOx rxn
between solid Fe (oxidized) and molecular oxygen
(reduced) water. The rxn causes a pit to form
that will eventually go all the way though the
metal. Some metals, such as Al or Ag, form a
protective oxidized layer that prevents pitting.
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How to Prevent Rust

• Coat the metal with paint or lacquer to seal out
  oxygen
• Use a sacrifice metal
• This implies a metal that oxidizes (rusts) easier
  than the metal you which to protect, is allowed
  to rust such that the electrons it produces are
  fed into the protected metal (see next slide)
• Apply an external D.C. voltage
• This is essentially the same as the previous
  method, but instead of a sacrifice metal a
  battery is used

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One Type of Cathodic Protection
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Tarnishing

• Some metals, such as aluminum and silver, form a
  thin layer of oxide (rust), which protects the
  metal from corroding or pitting all the way
  through
• Other metals, such as iron, have no such
  protection and rusts all the way through to the
  other side
• Sorry this is the last time we will use the
  slide notes this year!!!

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