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  • 1. CH 23 Electrochemistry 23.1 Electrochemical cells
  • 2. Types of electrochemical cells
    • Galvanic or Voltaic
    • The ‘spontaneous’ reaction.
    • Produces electrical energy.
    • Electrolytic
    • Non-spontaneous reaction.
    • Requires electrical energy to occur.
    • For reversible cells, the galvanic reaction can occur spontaneously and then be reversed electrolytically - rechargeable batteries.
  • 3. Voltaic cells
    • There are two general ways to conduct an oxidation-reduction reaction
    • Mixing oxidant and reductant together
    • Cu 2+ + Zn (s) Cu (s) + Zn 2+
    • This approach does not allow for control of the reaction.
  • 4. Voltaic cells
    • Electrochemical cells
    • Each half reaction is put in a separate ‘half cell.’ They can then be connected electrically.
    • This permits better control over the system.
  • 5. Spontaneous Reactions
    • Will occur if the anode metal is above the cathode metal in the Activity Series chart (pg 678)
  • 6. Voltaic Cell
    • Allessandro Volta (1745-1827) invented the first electrochemical cell- this type was called the voltaic cell.
    • It is a spontaneous reaction- he layered Cu and Zn plates, separated by cardboard: Cu plate had reduction occur, Zn plate had oxidation occur
  • 7. Voltaic Cell
    • A half-cell is one part of the voltaic cell where either oxidation or reduction occurs. A half cell consists of a metal strip immersed in a solution of it’s ions
  • 8. Voltaic cells Cu 2+ + Zn (s) Cu (s) + Zn 2+ Zn Cu Cu 2+ Zn 2+ e - e - Electrons are transferred from one half-cell to the other using an external metal conductor.
  • 9. Voltaic cells e - e - To complete the circuit, a salt bridge is used salt bridge
  • 10. Voltaic cells
    • Salt bridge Allows ion migration in solution but prevents extensive mixing of electrolytes.
    • It can be a simple porous disk or a gel saturated with a non-interfering, strong electrolyte like KCl .
    KCl Cl - K + Cl - is released to Zn side as Zn is converted to Zn 2+ K + is released as Cu 2+ is converted to Cu
  • 11. Voltaic cells For our example, we have zinc ion being produced. This is an oxidation so: The electrode is - the anode - is positive (+). “ AN OX” Zn Zn 2+ + 2e -
  • 12. Voltaic cells For our other half cell, we have copper metal being produced. This is a reduction so: The electrode is - the cathode - is negative (-) “ RED CAT” Cu 2+ + 2e - Cu
  • 13. Voltaic Cell
  • 14. Cell diagrams
    • Rather than drawing an entire cell, a type of shorthand can be used.
    • For our copper - zinc cell, it would be:
    • Zn | Zn 2+ (1M) || Cu 2+ (1M) |
    • The anode is always on the left.
    • | = boundaries between phases
    • || = salt bridge
    • Other conditions like concentration are listed just after each species.
  • 15. Dry Cell
    • Voltaic cell where the electrolyte is a paste- not a solution
    • Example: flashlight battery ( pg 681)
    • Not a true battery
    • Outer Zn case is anode (oxidation)
    • Carbon (graphite core) rod in center is cathode- but actually reduction occurs w/
    • MnO 2 found in paste
    • Salt bridge is not needed because of paste prevent cell contents from mixing
    • Alkaline batteries use KOH in paste and this makes it last longer and keeps voltage up
  • 16. Lead Storage Battery
    • A battery is a group of cells connected together
    • A car battery is 6 cells producing 2V each for a total of 12 V
    • The cathode is lead(IV) oxide and the anode is Pb. Dilute sulfuric acid is the electrolyte
    • Overall reaction is:
    • Pb (s) + PbO 2(s) + 2H 2 SO 4(aq) -----2PbSO 4(s) + 2 H 2 O (l)
    • Now you write the half reactions that occur at each electrode!!
  • 17. Lead Battery
    • Car battery’s are recharged when the car runs- the reaction occurs in reverse- but this reverse reaction is nonspontaneous and so the car’s generator supplies the energy to drive the reaction.
    • Eventually the battery dies- electrodes lose so much PbSO 4 which can fall to the bottom of the battery
  • 18. Fuel Cell
    • Idea here is to have a renewable electrode so electrodes don’t wear out
    • A fuel is used for the oxidation
    • Simplest is the Hydrogen-oxygen fuel cell- Oxygen is fiels for cathode (reduction), and hydrogen is fuel for anode (oxidation)
    • Overall reaction: 2H 2(g) + O 2(g) —2H 2 O (l)
  • 19. Fuel Cell
    • You write the anode and cathode half-cell reactions.
    • Advantage: cheap fuel, only “pollutant”- water which is drinkable
    • Used in spacecraft and some military applications- some cars; expensive and takes room.