Ch 23sec1


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Ch 23sec1

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